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Chapter 2

Molecules, Ions, and Chemical Formulas

Chapter 1 "Introduction to Chemistry" introduced some of the fundamentalconcepts of chemistry, with particular attention to the basic properties of atomsand elements. These entities are the building blocks of all substances we encounter,yet most common substances do not consist of only pure elements or individualatoms. Instead, nearly all substances are chemical compounds or mixtures ofchemical compounds. Although there are only about 115 elements (of which about86 occur naturally), millions of chemical compounds are known, with a tremendousrange of physical and chemical properties. Consequently, the emphasis of modernchemistry (and this text) is on understanding the relationship between thestructures and properties of chemical compounds.

Petroleum refining. Using chemicals, catalysts, heat, and pressure, a petroleum refinery will separate, combine,and rearrange the structure and bonding patterns of the basic carbon-hydrogen molecules found in crude oil. Thefinal products include gasoline, paraffin, diesel fuel, lubricants, and bitumen.

123

In this chapter, you will learn how to describe the composition of chemicalcompounds. We introduce you to chemical nomenclature—the language ofchemistry—that will enable you to recognize and name the most common kinds ofcompounds. An understanding of chemical nomenclature not only is essential foryour study of chemistry but also has other benefits—for example, it helps youunderstand the labels on products found in the supermarket and the pharmacy.You will also be better equipped to understand many of the importantenvironmental and medical issues that face society. By the end of this chapter, youwill be able to describe what happens chemically when a doctor prepares a cast tostabilize a broken bone, and you will know the composition of common substancessuch as laundry bleach, the active ingredient in baking powder, and the foul-smelling compound responsible for the odor of spoiled fish. Finally, you will be ableto explain the chemical differences among different grades of gasoline.

Chapter 2 Molecules, Ions, and Chemical Formulas

124

2.1 Chemical Compounds

LEARNING OBJECTIVE

1. To understand the differences between covalent and ionic bonding.

The atoms in all substances that contain more than one atom are held together byelectrostatic interactions1—interactions between electrically charged particlessuch as protons and electrons. Electrostatic attraction2 between oppositelycharged species (positive and negative) results in a force that causes them to movetoward each other, like the attraction between opposite poles of two magnets. Incontrast, electrostatic repulsion3 between two species with the same charge(either both positive or both negative) results in a force that causes them to repeleach other, as do the same poles of two magnets. Atoms form chemical compoundswhen the attractive electrostatic interactions between them are stronger than therepulsive interactions. Collectively, we refer to the attractive interactions betweenatoms as chemical bonds4.

Chemical bonds are generally divided into two fundamentally different kinds: ionicand covalent. In reality, however, the bonds in most substances are neither purelyionic nor purely covalent, but they are closer to one of these extremes. Althoughpurely ionic and purely covalent bonds represent extreme cases that are seldomencountered in anything but very simple substances, a brief discussion of these twoextremes helps us understand why substances that have different kinds of chemicalbonds have very different properties. Ionic compounds5 consist of positively andnegatively charged ions held together by strong electrostatic forces, whereascovalent compounds6 generally consist of molecules7, which are groups of atomsin which one or more pairs of electrons are shared between bonded atoms. In acovalent bond8, the atoms are held together by the electrostatic attraction betweenthe positively charged nuclei of the bonded atoms and the negatively chargedelectrons they share. We begin our discussion of structures and formulas bydescribing covalent compounds. The energetic factors involved in bond formationare described in more quantitative detail in Chapter 8 "Ionic versus CovalentBonding".

1. An interaction betweenelectrically charged particlessuch as protons and electrons.

2. An electrostatic interactionbetween oppositely chargedspecies (positive and negative)that results in a force thatcauses them to move towardeach other.

3. An electrostatic interactionbetween two species that havethe same charge (both positiveor both negative) that resultsin a force that causes them torepel each other.

4. An attractive interactionbetween atoms that holds themtogether in compounds.

5. A compound consisting ofpositively charged ions(cations) and negativelycharged ions (anions) heldtogether by strongelectrostatic forces.

6. A compound that consists ofdiscrete molecules.

7. A group of atoms in which oneor more pairs of electrons areshared between bonded atoms.

8. The electrostatic attractionbetween the positively chargednuclei of the bonded atoms andthe negatively chargedelectrons they share.


125

Note the Pattern

Ionic compounds consist of ions of opposite charges held together by strongelectrostatic forces, whereas pairs of electrons are shared between bondedatoms in covalent compounds.

Covalent Molecules and Compounds

Just as an atom is the simplest unit that has the fundamental chemical properties ofan element, a molecule is the simplest unit that has the fundamental chemicalproperties of a covalent compound. Some pure elements exist as covalentmolecules. Hydrogen, nitrogen, oxygen, and the halogens occur naturally as thediatomic (“two atoms”) molecules H2, N2, O2, F2, Cl2, Br2, and I2 (part (a) in Figure 2.1

"Elements That Exist as Covalent Molecules"). Similarly, a few pure elements arepolyatomic9 (“many atoms”) molecules, such as elemental phosphorus and sulfur,which occur as P4 and S8 (part (b) in Figure 2.1 "Elements That Exist as Covalent

Molecules").

Each covalent compound is represented by a molecular formula10, which gives theatomic symbol for each component element, in a prescribed order, accompanied bya subscript indicating the number of atoms of that element in the molecule. Thesubscript is written only if the number of atoms is greater than 1. For example,water, with two hydrogen atoms and one oxygen atom per molecule, is written asH2O. Similarly, carbon dioxide, which contains one carbon atom and two oxygen

atoms in each molecule, is written as CO2.

Figure 2.1 Elements That Exist as Covalent Molecules

9. Molecules that contain morethan two atoms.

10. A representation of a covalentcompound that consists of theatomic symbol for eachcomponent element (in aprescribed order) accompaniedby a subscript indicating thenumber of atoms of thatelement in the molecule. Thesubscript is written only if thenumber is greater than 1.


2.1 Chemical Compounds 126

(a) Several elements naturally exist as diatomic molecules, in which two atoms (E) are joined by one or morecovalent bonds to form a molecule with the general formula E2. (b) A few elements naturally exist as polyatomic

molecules, which contain more than two atoms. For example, phosphorus exists as P4 tetrahedra—regular polyhedra

with four triangular sides—with a phosphorus atom at each vertex. Elemental sulfur consists of a puckered ring ofeight sulfur atoms connected by single bonds. Selenium is not shown due to the complexity of its structure.

Covalent compounds that contain predominantly carbon and hydrogen are calledorganic compounds11. The convention for representing the formulas of organiccompounds is to write carbon first, followed by hydrogen and then any otherelements in alphabetical order (e.g., CH4O is methyl alcohol, a fuel). Compounds

that consist primarily of elements other than carbon and hydrogen are calledinorganic compounds12; they include both covalent and ionic compounds. Ininorganic compounds, the component elements are listed beginning with the onefarthest to the left in the periodic table (see Chapter 32 "Appendix H: Periodic Tableof Elements"), such as we see in CO2 or SF6. Those in the same group are listed

beginning with the lower element and working up, as in ClF. By convention,however, when an inorganic compound contains both hydrogen and an elementfrom groups 13–15, the hydrogen is usually listed last in the formula. Examples areammonia (NH3) and silane (SiH4). Compounds such as water, whose compositions

were established long before this convention was adopted, are always written withhydrogen first: Water is always written as H2O, not OH2. The conventions for

inorganic acids, such as hydrochloric acid (HCl) and sulfuric acid (H2SO4), are

described in Section 2.5 "Acids and Bases".

Note the Pattern

For organic compounds: write C first, then H, and then the other elements inalphabetical order. For molecular inorganic compounds: start with the elementat far left in the periodic table; list elements in same group beginning with thelower element and working up.

11. A covalent compound thatcontains predominantly carbonand hydrogen.

12. An ionic or covalent compoundthat consists primarily ofelements other than carbonand hydrogen.



averill_1.0-ch96appH#averill_1.0-ch96appH

averill_1.0-ch96appH#averill_1.0-ch96appH

EXAMPLE 1

Write the molecular formula of each compound.

a. The phosphorus-sulfur compound that is responsible for the ignition ofso-called strike anywhere matches has 4 phosphorus atoms and 3 sulfuratoms per molecule.

b. Ethyl alcohol, the alcohol of alcoholic beverages, has 1 oxygen atom, 2carbon atoms, and 6 hydrogen atoms per molecule.

c. Freon-11, once widely used in automobile air conditioners andimplicated in damage to the ozone layer, has 1 carbon atom, 3 chlorineatoms, and 1 fluorine atom per molecule.

Given: identity of elements present and number of atoms of each

Asked for: molecular formula

Strategy:

A Identify the symbol for each element in the molecule. Then identify thesubstance as either an organic compound or an inorganic compound.

B If the substance is an organic compound, arrange the elements in orderbeginning with carbon and hydrogen and then list the other elementsalphabetically. If it is an inorganic compound, list the elements beginningwith the one farthest left in the periodic table. List elements in the samegroup starting with the lower element and working up.

C From the information given, add a subscript for each kind of atom to writethe molecular formula.

Solution:

a. A The molecule has 4 phosphorus atoms and 3 sulfur atoms. Because thecompound does not contain mostly carbon and hydrogen, it is inorganic.B Phosphorus is in group 15, and sulfur is in group 16. Becausephosphorus is to the left of sulfur, it is written first. C Writing thenumber of each kind of atom as a right-hand subscript gives P4S3 as themolecular formula.

b. A Ethyl alcohol contains predominantly carbon and hydrogen, so it is anorganic compound. B The formula for an organic compound is written



with the number of carbon atoms first, the number of hydrogen atomsnext, and the other atoms in alphabetical order: CHO. C Addingsubscripts gives the molecular formula C2H6O.

c. A Freon-11 contains carbon, chlorine, and fluorine. It can beviewed as either an inorganic compound or an organiccompound (in which fluorine has replaced hydrogen). Theformula for Freon-11 can therefore be written using either of thetwo conventions.

B According to the convention for inorganic compounds, carbonis written first because it is farther left in the periodic table.Fluorine and chlorine are in the same group, so they are listedbeginning with the lower element and working up: CClF. Addingsubscripts gives the molecular formula CCl3F.

C We obtain the same formula for Freon-11 using the conventionfor organic compounds. The number of carbon atoms is writtenfirst, followed by the number of hydrogen atoms (zero) and thenthe other elements in alphabetical order, also giving CCl3F.

Exercise

Write the molecular formula for each compound.

a. Nitrous oxide, also called “laughing gas,” has 2 nitrogen atoms and 1oxygen atom per molecule. Nitrous oxide is used as a mild anesthetic forminor surgery and as the propellant in cans of whipped cream.

b. Sucrose, also known as cane sugar, has 12 carbon atoms, 11 oxygenatoms, and 22 hydrogen atoms.

c. Sulfur hexafluoride, a gas used to pressurize “unpressurized” tennisballs and as a coolant in nuclear reactors, has 6 fluorine atoms and 1sulfur atom per molecule.

Answer:

a. N2Ob. C12H22O11

c. SF6



Representations of Molecular Structures

Molecular formulas give only the elemental composition of molecules. In contrast,structural formulas13 show which atoms are bonded to one another and, in somecases, the approximate arrangement of the atoms in space. Knowing the structuralformula of a compound enables chemists to create a three-dimensional model,which provides information about how that compound will behave physically andchemically.

The structural formula for H2 can be drawn as H–H and that for I2 as I–I, where the

line indicates a single pair of shared electrons, a single bond14. Two pairs ofelectrons are shared in a double bond15, which is indicated by two lines— forexample, O2 is O=O. Three electron pairs are shared in a triple bond16, which is

indicated by three lines—for example, N2 is N≡N (see Figure 2.2 "Molecules That

Contain Single, Double, and Triple Bonds"). Carbon is unique in the extent to whichit forms single, double, and triple bonds to itself and other elements. The number ofbonds formed by an atom in its covalent compounds is not arbitrary. As you willlearn in Chapter 8 "Ionic versus Covalent Bonding", hydrogen, oxygen, nitrogen,and carbon have a very strong tendency to form substances in which they have one,two, three, and four bonds to other atoms, respectively (Table 2.1 "The Number ofBonds That Selected Atoms Commonly Form to Other Atoms").

Figure 2.2 Molecules That Contain Single, Double, and Triple Bonds

13. A representation of a moleculethat shows which atoms arebonded to one another and, insome cases, the approximatearrangement of atoms in space.

14. A chemical bond formed whentwo atoms share a single pairof electrons.

15. A chemical bond formed whentwo atoms share two pairs ofelectrons.

16. A chemical bond formed whentwo atoms share three pairs ofelectrons.



Hydrogen (H2) has a single bond between atoms. Oxygen (O2) has a double bond between atoms, indicated by two

lines (=). Nitrogen (N2) has a triple bond between atoms, indicated by three lines (≡). Each bond represents an

electron pair.

Table 2.1 The Number of Bonds That Selected Atoms Commonly Form to OtherAtoms

Atom Number of Bonds

H (group 1) 1

O (group 16) 2

N (group 15) 3

C (group 14) 4

The structural formula for water can be drawn as follows:

Because the latter approximates the experimentally determined shape of the watermolecule, it is more informative. Similarly, ammonia (NH3) and methane (CH4) are

often written as planar molecules:

As shown in Figure 2.3 "The Three-Dimensional Structures of Water, Ammonia, andMethane", however, the actual three-dimensional structure of NH3 looks like a

pyramid with a triangular base of three hydrogen atoms. The structure of CH4, with

four hydrogen atoms arranged around a central carbon atom as shown in Figure 2.3"The Three-Dimensional Structures of Water, Ammonia, and Methane", istetrahedral. That is, the hydrogen atoms are positioned at every other vertex of acube. Many compounds—carbon compounds, in particular—have four bonded atomsarranged around a central atom to form a tetrahedron.



CH4. Methane has a three-

dimensional, tetrahedralstructure.

Figure 2.3 The Three-Dimensional Structures of Water, Ammonia, and Methane

(a) Water is a V-shaped molecule, in which all three atoms lie in a plane. (b) In contrast, ammonia has a pyramidalstructure, in which the three hydrogen atoms form the base of the pyramid and the nitrogen atom is at the vertex.(c) The four hydrogen atoms of methane form a tetrahedron; the carbon atom lies in the center.

Figure 2.1 "Elements That Exist as Covalent Molecules",Figure 2.2 "Molecules That Contain Single, Double, andTriple Bonds", and Figure 2.3 "The Three-DimensionalStructures of Water, Ammonia, and Methane" illustratedifferent ways to represent the structures of molecules.It should be clear that there is no single “best” way todraw the structure of a molecule; the method you usedepends on which aspect of the structure you want toemphasize and how much time and effort you want tospend. Figure 2.4 "Different Ways of Representing theStructure of a Molecule" shows some of the differentways to portray the structure of a slightly more complexmolecule: methanol. These representations differgreatly in their information content. For example, themolecular formula for methanol (part (a) in Figure 2.4"Different Ways of Representing the Structure of aMolecule") gives only the number of each kind of atom;writing methanol as CH4O tells nothing about its

structure. In contrast, the structural formula (part (b) in Figure 2.4 "Different Waysof Representing the Structure of a Molecule") indicates how the atoms areconnected, but it makes methanol look as if it is planar (which it is not). Both theball-and-stick model (part (c) in Figure 2.4 "Different Ways of Representing theStructure of a Molecule") and the perspective drawing (part (d) in Figure 2.4"Different Ways of Representing the Structure of a Molecule") show the three-



dimensional structure of the molecule. The latter (also called a wedge-and-dashrepresentation) is the easiest way to sketch the structure of a molecule in threedimensions. It shows which atoms are above and below the plane of the paper byusing wedges and dashes, respectively; the central atom is always assumed to be inthe plane of the paper. The space-filling model (part (e) in Figure 2.4 "DifferentWays of Representing the Structure of a Molecule") illustrates the approximaterelative sizes of the atoms in the molecule, but it does not show the bonds betweenthe atoms. Also, in a space-filling model, atoms at the “front” of the molecule mayobscure atoms at the “back.”

Figure 2.4 Different Ways of Representing the Structure of a Molecule

(a) The molecular formula for methanol gives only the number of each kind of atom present. (b) The structuralformula shows which atoms are connected. (c) The ball-and-stick model shows the atoms as spheres and the bondsas sticks. (d) A perspective drawing (also called a wedge-and-dash representation) attempts to show the three-dimensional structure of the molecule. (e) The space-filling model shows the atoms in the molecule but not thebonds. (f) The condensed structural formula is by far the easiest and most common way to represent a molecule.

Although a structural formula, a ball-and-stick model, a perspective drawing, and aspace-filling model provide a significant amount of information about the structureof a molecule, each requires time and effort. Consequently, chemists often use acondensed structural formula (part (f) in Figure 2.4 "Different Ways of Representingthe Structure of a Molecule"), which omits the lines representing bonds betweenatoms and simply lists the atoms bonded to a given atom next to it. Multiple groupsattached to the same atom are shown in parentheses, followed by a subscript thatindicates the number of such groups. For example, the condensed structuralformula for methanol is CH3OH, which tells us that the molecule contains a CH3 unit

that looks like a fragment of methane (CH4). Methanol can therefore be viewed

either as a methane molecule in which one hydrogen atom has been replaced by an–OH group or as a water molecule in which one hydrogen atom has been replacedby a –CH3 fragment. Because of their ease of use and information content, we use

condensed structural formulas for molecules throughout this text. Ball-and-stickmodels are used when needed to illustrate the three-dimensional structure ofmolecules, and space-filling models are used only when it is necessary to visualizethe relative sizes of atoms or molecules to understand an important point.



EXAMPLE 2

Write the molecular formula for each compound. The condensed structuralformula is given.

a. Sulfur monochloride (also called disulfur dichloride) is a vile-smelling,corrosive yellow liquid used in the production of synthetic rubber. Itscondensed structural formula is ClSSCl.

b. Ethylene glycol is the major ingredient in antifreeze. Its condensedstructural formula is HOCH2CH2OH.

c. Trimethylamine is one of the substances responsible for the smell ofspoiled fish. Its condensed structural formula is (CH3)3N.

Given: condensed structural formula

Asked for: molecular formula

Strategy:

A Identify every element in the condensed structural formula and thendetermine whether the compound is organic or inorganic.

B As appropriate, use either organic or inorganic convention to list theelements. Then add appropriate subscripts to indicate the number of atomsof each element present in the molecular formula.

Solution:

The molecular formula lists the elements in the molecule and the number ofatoms of each.

a. A Each molecule of sulfur monochloride has two sulfur atoms and twochlorine atoms. Because it does not contain mostly carbon andhydrogen, it is an inorganic compound. B Sulfur lies to the left ofchlorine in the periodic table, so it is written first in the formula. Addingsubscripts gives the molecular formula S2Cl2.

b. A Counting the atoms in ethylene glycol, we get six hydrogen atoms, twocarbon atoms, and two oxygen atoms per molecule. The compoundconsists mostly of carbon and hydrogen atoms, so it is organic. B As withall organic compounds, C and H are written first in the molecular



formula. Adding appropriate subscripts gives the molecular formulaC2H6O2.

c. A The condensed structural formula shows that trimethylaminecontains three CH3 units, so we have one nitrogen atom, three carbonatoms, and nine hydrogen atoms per molecule. Because trimethylaminecontains mostly carbon and hydrogen, it is an organic compound. BAccording to the convention for organic compounds, C and H arewritten first, giving the molecular formula C3H9N.

Exercise

Write the molecular formula for each molecule.

a. Chloroform, which was one of the first anesthetics and was used inmany cough syrups until recently, contains one carbon atom, onehydrogen atom, and three chlorine atoms. Its condensed structuralformula is CHCl3.

b. Hydrazine is used as a propellant in the attitude jets of the space shuttle.Its condensed structural formula is H2NNH2.

c. Putrescine is a pungent-smelling compound first isolated from extractsof rotting meat. Its condensed structural formula isH2NCH2CH2CH2CH2NH2. This is often written as H2N(CH2)4NH2 toindicate that there are four CH2 fragments linked together.



Answer:

a. CHCl3b. N2H4

c. C4H12N2

Ionic Compounds

The substances described in the preceding discussion are composed of moleculesthat are electrically neutral; that is, the number of positively charged protons in thenucleus is equal to the number of negatively charged electrons. In contrast, ions areatoms or assemblies of atoms that have a net electrical charge. Ions that containfewer electrons than protons have a net positive charge and are called cations17.Conversely, ions that contain more electrons than protons have a net negativecharge and are called anions18. Ionic compounds contain both cations and anions in aratio that results in no net electrical charge.

Note the Pattern

Ionic compounds contain both cations and anions in a ratio that results in zeroelectrical charge.

In covalent compounds, electrons are shared between bonded atoms and aresimultaneously attracted to more than one nucleus. In contrast, ionic compoundscontain cations and anions rather than discrete neutral molecules. Ionic compoundsare held together by the attractive electrostatic interactions between cations andanions. In an ionic compound, the cations and anions are arranged in space to forman extended three-dimensional array that maximizes the number of attractiveelectrostatic interactions and minimizes the number of repulsive electrostaticinteractions (Figure 2.5 "Covalent and Ionic Bonding"). As shown in Equation 2.1,the electrostatic energy of the interaction between two charged particles isproportional to the product of the charges on the particles and inverselyproportional to the distance between them:

Equation 2.1

electrostatic energy ∝Q1Q2

r

17. An ion that has fewer electronsthan protons, resulting in a netpositive charge.

18. An ion that has fewer protonsthan electrons, resulting in anet negative charge.



where Q1 and Q2 are the electrical charges on particles 1 and 2, and r is the distance

between them. When Q1 and Q2 are both positive, corresponding to the charges on

cations, the cations repel each other and the electrostatic energy is positive. WhenQ1 and Q2 are both negative, corresponding to the charges on anions, the anions

repel each other and the electrostatic energy is again positive. The electrostaticenergy is negative only when the charges have opposite signs; that is, positivelycharged species are attracted to negatively charged species and vice versa. Asshown in Figure 2.6 "The Effect of Charge and Distance on the Strength ofElectrostatic Interactions", the strength of the interaction is proportional to themagnitude of the charges and decreases as the distance between the particlesincreases. We will return to these energetic factors in Chapter 8 "Ionic versusCovalent Bonding", where they are described in greater quantitative detail.

Note the Pattern

If the electrostatic energy is positive, the particles repel each other; if theelectrostatic energy is negative, the particles are attracted to each other.

Figure 2.5 Covalent and Ionic Bonding

(a) In molecular hydrogen (H2), two hydrogen atoms share two electrons to form a covalent bond. (b) The ionic

compound NaCl forms when electrons from sodium atoms are transferred to chlorine atoms. The resulting Na+ andCl− ions form a three-dimensional solid that is held together by attractive electrostatic interactions.



Figure 2.6 The Effect of Charge and Distance on the Strength of Electrostatic Interactions

As the charge on ions increases or the distance between ions decreases, so does the strength of the attractive (−…+)or repulsive (−…− or +…+) interactions. The strength of these interactions is represented by the thickness of thearrows.

One example of an ionic compound is sodium chloride (NaCl; Figure 2.7 "SodiumChloride: an Ionic Solid"), formed from sodium and chlorine. In forming chemicalcompounds, many elements have a tendency to gain or lose enough electrons toattain the same number of electrons as the noble gas closest to them in the periodictable. When sodium and chlorine come into contact, each sodium atom gives up anelectron to become a Na+ ion, with 11 protons in its nucleus but only 10 electrons(like neon), and each chlorine atom gains an electron to become a Cl− ion, with 17protons in its nucleus and 18 electrons (like argon), as shown in part (b) in Figure2.5 "Covalent and Ionic Bonding". Solid sodium chloride contains equal numbers ofcations (Na+) and anions (Cl−), thus maintaining electrical neutrality. Each Na+ ion issurrounded by 6 Cl− ions, and each Cl− ion is surrounded by 6 Na+ ions. Because ofthe large number of attractive Na+Cl− interactions, the total attractive electrostaticenergy in NaCl is great.



Figure 2.7 Sodium Chloride: an Ionic Solid

The planes of an NaCl crystal reflect the regular three-dimensional arrangement of its Na+ (purple) and Cl− (green)ions.

Consistent with a tendency to have the same number of electrons as the nearestnoble gas, when forming ions, elements in groups 1, 2, and 3 tend to lose one, two,and three electrons, respectively, to form cations, such as Na+ and Mg2+. They thenhave the same number of electrons as the nearest noble gas: neon. Similarly, K+,Ca2+, and Sc3+ have 18 electrons each, like the nearest noble gas: argon. In addition,the elements in group 13 lose three electrons to form cations, such as Al3+, againattaining the same number of electrons as the noble gas closest to them in theperiodic table. Because the lanthanides and actinides formally belong to group 3,the most common ion formed by these elements is M3+, where M represents themetal. Conversely, elements in groups 17, 16, and 15 often react to gain one, two,and three electrons, respectively, to form ions such as Cl−, S2−, and P3−. Ions such asthese, which contain only a single atom, are called monatomic ions19. You canpredict the charges of most monatomic ions derived from the main group elementsby simply looking at the periodic table and counting how many columns an elementlies from the extreme left or right. For example, you can predict that barium (ingroup 2) will form Ba2+ to have the same number of electrons as its nearest noble19. An ion with only a single atom.



gas, xenon, that oxygen (in group 16) will form O2− to have the same number ofelectrons as neon, and cesium (in group 1) will form Cs+ to also have the samenumber of electrons as xenon. Note that this method does not usually work formost of the transition metals, as you will learn in Section 2.3 "Naming IonicCompounds". Some common monatomic ions are in Table 2.2 "Some CommonMonatomic Ions and Their Names".

Note the Pattern

Elements in groups 1, 2, and 3 tend to form 1+, 2+, and 3+ ions, respectively;elements in groups 15, 16, and 17 tend to form 3−, 2−, and 1− ions, respectively.

Table 2.2 Some Common Monatomic Ions and Their Names

Group 1 Group 2 Group 3 Group 13 Group 15 Group 16 Group 17

Li+

lithium

Be2+

beryllium

N3−

nitride

(azide)

O2−

oxide

F−

fluoride

Na+

sodium

Mg2+

magnesium

Al3+

aluminum

P3−

phosphide

S2−

sulfide

Cl−

chloride

K+

potassium

Ca2+

calcium

Sc3+

scandium

Ga3+

gallium

As3−

arsenide

Se2−

selenide

Br−

bromide



Group 1 Group 2 Group 3 Group 13 Group 15 Group 16 Group 17

Rb+

rubidium

Sr2+

strontium

Y3+

yttrium

In3+

indium

Te2−

telluride

I−

iodide

Cs+

cesium

Ba2+

barium

La3+

lanthanum



EXAMPLE 3

Predict the charge on the most common monatomic ion formed by eachelement.

a. aluminum, used in the quantum logic clock, the world’s most preciseclock

b. selenium, used to make ruby-colored glassc. yttrium, used to make high-performance spark plugs

Given: element

Asked for: ionic charge

Strategy:

A Identify the group in the periodic table to which the element belongs.Based on its location in the periodic table, decide whether the element is ametal, which tends to lose electrons; a nonmetal, which tends to gainelectrons; or a semimetal, which can do either.

B After locating the noble gas that is closest to the element, determine thenumber of electrons the element must gain or lose to have the same numberof electrons as the nearest noble gas.

Solution:

a. A Aluminum is a metal in group 13; consequently, it will tend to loseelectrons. B The nearest noble gas to aluminum is neon. Aluminum willlose three electrons to form the Al3+ ion, which has the same number ofelectrons as neon.

b. A Selenium is a nonmetal in group 16, so it will tend to gain electrons. BThe nearest noble gas is krypton, so we predict that selenium will gaintwo electrons to form the Se2− ion, which has the same number ofelectrons as krypton.

c. A Yttrium is in group 3, and elements in this group are metals that tendto lose electrons. B The nearest noble gas to yttrium is krypton, soyttrium is predicted to lose three electrons to form Y3+, which has thesame number of electrons as krypton.

Exercise



Predict the charge on the most common monatomic ion formed by eachelement.

a. calcium, used to prevent osteoporosisb. iodine, required for the synthesis of thyroid hormonesc. zirconium, widely used in nuclear reactors

Answer:

a. Ca2+

b. I−

c. Zr4+

Physical Properties of Ionic and Covalent Compounds

In general, ionic and covalent compounds have different physical properties. Ioniccompounds usually form hard crystalline solids that melt at rather hightemperatures and are very resistant to evaporation. These properties stem from thecharacteristic internal structure of an ionic solid, illustrated schematically in part(a) in Figure 2.8 "Interactions in Ionic and Covalent Solids", which shows the three-dimensional array of alternating positive and negative ions held together by strongelectrostatic attractions. In contrast, as shown in part (b) in Figure 2.8 "Interactionsin Ionic and Covalent Solids", most covalent compounds consist of discretemolecules held together by comparatively weak intermolecular forces (the forcesbetween molecules), even though the atoms within each molecule are held togetherby strong intramolecular covalent bonds (the forces within the molecule). Covalentsubstances can be gases, liquids, or solids at room temperature and pressure,depending on the strength of the intermolecular interactions. Covalent molecularsolids tend to form soft crystals that melt at rather low temperatures and evaporaterelatively easily.Some covalent substances, however, are not molecular but consistof infinite three-dimensional arrays of covalently bonded atoms and include someof the hardest materials known, such as diamond. This topic will be addressed inChapter 12 "Solids". The covalent bonds that hold the atoms together in themolecules are unaffected when covalent substances melt or evaporate, so a liquid orvapor of discrete, independent molecules is formed. For example, at roomtemperature, methane, the major constituent of natural gas, is a gas that iscomposed of discrete CH4 molecules. A comparison of the different physical

properties of ionic compounds and covalent molecular substances is given in Table2.3 "The Physical Properties of Typical Ionic Compounds and Covalent MolecularSubstances".



Table 2.3 The Physical Properties of Typical Ionic Compounds and CovalentMolecular Substances

Ionic Compounds Covalent Molecular Substances

hard solids gases, liquids, or soft solids

high melting points low melting points

nonvolatile volatile

Figure 2.8 Interactions in Ionic and Covalent Solids

(a) The positively and negatively charged ions in an ionic solid such as sodium chloride (NaCl) are held together bystrong electrostatic interactions. (b) In this representation of the packing of methane (CH4) molecules in solid

methane, a prototypical molecular solid, the methane molecules are held together in the solid only by relativelyweak intermolecular forces, even though the atoms within each methane molecule are held together by strongcovalent bonds.



Summary

The atoms in chemical compounds are held together by attractive electrostaticinteractions known as chemical bonds. Ionic compounds contain positivelyand negatively charged ions in a ratio that results in an overall charge of zero.The ions are held together in a regular spatial arrangement by electrostaticforces. Most covalent compounds consist of molecules, groups of atoms inwhich one or more pairs of electrons are shared by at least two atoms to form acovalent bond. The atoms in molecules are held together by the electrostaticattraction between the positively charged nuclei of the bonded atoms and thenegatively charged electrons shared by the nuclei. The molecular formula of acovalent compound gives the types and numbers of atoms present. Compoundsthat contain predominantly carbon and hydrogen are called organiccompounds, whereas compounds that consist primarily of elements other thancarbon and hydrogen are inorganic compounds. Diatomic molecules contain twoatoms, and polyatomic molecules contain more than two. A structural formulaindicates the composition and approximate structure and shape of a molecule.Single bonds, double bonds, and triple bonds are covalent bonds in whichone, two, and three pairs of electrons, respectively, are shared between twobonded atoms. Atoms or groups of atoms that possess a net electrical chargeare called ions; they can have either a positive charge (cations) or a negativecharge (anions). Ions can consist of one atom (monatomic ions) or several(polyatomic ions). The charges on monatomic ions of most main group elementscan be predicted from the location of the element in the periodic table. Ioniccompounds usually form hard crystalline solids with high melting points.Covalent molecular compounds, in contrast, consist of discrete molecules heldtogether by weak intermolecular forces and can be gases, liquids, or solids atroom temperature and pressure.

KEY TAKEAWAY

• There are two fundamentally different kinds of chemical bonds(covalent and ionic) that cause substances to have very differentproperties.



CONCEPTUAL PROBLEMS

1. Ionic and covalent compounds are held together by electrostatic attractionsbetween oppositely charged particles. Describe the differences in the nature ofthe attractions in ionic and covalent compounds. Which class of compoundscontains pairs of electrons shared between bonded atoms?

2. Which contains fewer electrons than the neutral atom—the correspondingcation or the anion?

3. What is the difference between an organic compound and an inorganiccompound?

4. What is the advantage of writing a structural formula as a condensed formula?

5. The majority of elements that exist as diatomic molecules are found in onegroup of the periodic table. Identify the group.

6. Discuss the differences between covalent and ionic compounds with regard to

a. the forces that hold the atoms together.b. melting points.c. physical states at room temperature and pressure.

7. Why do covalent compounds generally tend to have lower melting points thanionic compounds?

ANSWER

7. Covalent compounds generally melt at lower temperatures than ioniccompounds because the intermolecular interactions that hold the moleculestogether in a molecular solid are weaker than the electrostatic attractions thathold oppositely charged ions together in an ionic solid.



NUMERICAL PROBLEMS

1. The structural formula for chloroform (CHCl3) was shown in Example 2. Basedon this information, draw the structural formula of dichloromethane (CH2Cl2).

2. What is the total number of electrons present in each ion?

a. F−

b. Rb+

c. Ce3+

d. Zr4+

e. Zn2+

f. Kr2+

g. B3+

3. What is the total number of electrons present in each ion?

a. Ca2+

b. Se2−

c. In3+

d. Sr2+

e. As3+

f. N3−

g. Tl+

4. Predict how many electrons are in each ion.

a. an oxygen ion with a −2 chargeb. a beryllium ion with a +2 chargec. a silver ion with a +1 charged. a selenium ion with a +4 chargee. an iron ion with a +2 chargef. a chlorine ion with a −1 charge

5. Predict how many electrons are in each ion.

a. a copper ion with a +2 chargeb. a molybdenum ion with a +4 chargec. an iodine ion with a −1 charged. a gallium ion with a +3 chargee. an ytterbium ion with a +3 chargef. a scandium ion with a +3 charge

6. Predict the charge on the most common monatomic ion formed by eachelement.



a. chlorineb. phosphorusc. scandiumd. magnesiume. arsenicf. oxygen

7. Predict the charge on the most common monatomic ion formed by eachelement.

a. sodiumb. seleniumc. bariumd. rubidiume. nitrogenf. aluminum

8. For each representation of a monatomic ion, identify the parent atom, writethe formula of the ion using an appropriate superscript, and indicate theperiod and group of the periodic table in which the element is found.

a. X49 2+

b. X11 –

c. X816 2–

9. For each representation of a monatomic ion, identify the parent atom, writethe formula of the ion using an appropriate superscript, and indicate theperiod and group of the periodic table in which the element is found.

a. X37 +

b. X919 –

c. X1327 3+



ANSWERS

5. a. 27b. 38c. 54d. 28e. 67f. 18

9. a. Li, Li+, 2nd period, group 1b. F, F–, 2nd period, group 17c. Al, Al3+, 3nd period, group 13



2.2 Chemical Formulas

LEARNING OBJECTIVE

1. To describe the composition of a chemical compound.

When chemists synthesize a new compound, they may not yet know its molecularor structural formula. In such cases, they usually begin by determining itsempirical formula20, the relative numbers of atoms of the elements in a compound,reduced to the smallest whole numbers. Because the empirical formula is based onexperimental measurements of the numbers of atoms in a sample of the compound,it shows only the ratios of the numbers of the elements present. The differencebetween empirical and molecular formulas can be illustrated with butane, a covalentcompound used as the fuel in disposable lighters. The molecular formula for butaneis C4H10. The ratio of carbon atoms to hydrogen atoms in butane is 4:10, which can

be reduced to 2:5. The empirical formula for butane is therefore C2H5. The formula

unit21 is the absolute grouping of atoms or ions represented by the empiricalformula of a compound, either ionic or covalent. Butane, for example, has theempirical formula C2H5, but it contains two C2H5 formula units, giving a molecular

formula of C4H10.

Because ionic compounds do not contain discrete molecules, empirical formulas areused to indicate their compositions. All compounds, whether ionic or covalent,must be electrically neutral. Consequently, the positive and negative charges in aformula unit must exactly cancel each other. If the cation and the anion havecharges of equal magnitude, such as Na+ and Cl−, then the compound must have a1:1 ratio of cations to anions, and the empirical formula must be NaCl. If the chargesare not the same magnitude, then a cation:anion ratio other than 1:1 is needed toproduce a neutral compound. In the case of Mg2+ and Cl−, for example, two Cl− ionsare needed to balance the two positive charges on each Mg2+ ion, giving anempirical formula of MgCl2. Similarly, the formula for the ionic compound that

contains Na+ and O2− ions is Na2O.

20. A formula for a compound thatconsists of the atomic symbolfor each component elementaccompanied by a subscriptindicating the relative numberof atoms of that element in thecompound, reduced to thesmallest whole numbers.

21. The absolute grouping of atomsor ions represented by theempirical formula.


150

Note the Pattern

Ionic compounds do not contain discrete molecules, so empirical formulas areused to indicate their compositions.

Binary Ionic Compounds

An ionic compound that contains only two elements, one present as a cation andone as an anion, is called a binary ionic compound22. One example is MgCl2, a

coagulant used in the preparation of tofu from soybeans. For binary ioniccompounds, the subscripts in the empirical formula can also be obtained bycrossing charges: use the absolute value of the charge on one ion as the subscriptfor the other ion. This method is shown schematically as follows:

Crossing charges. One method for obtaining subscripts in the empirical formula is by crossing charges.

When crossing charges, you will sometimes find it necessary to reduce thesubscripts to their simplest ratio to write the empirical formula. Consider, forexample, the compound formed by Mg2+ and O2−. Using the absolute values of thecharges on the ions as subscripts gives the formula Mg2O2:

This simplifies to its correct empirical formula MgO. The empirical formula has oneMg2+ ion and one O2− ion.

22. An ionic compound thatcontains only two elements,one present as a cation and oneas an anion.


2.2 Chemical Formulas 151

EXAMPLE 4

Write the empirical formula for the simplest binary ionic compound formedfrom each ion or element pair.

a. Ga3+ and As3−

b. Eu3+ and O2−

c. calcium and chlorine

Given: ions or elements

Asked for: empirical formula for binary ionic compound

Strategy:

A If not given, determine the ionic charges based on the location of theelements in the periodic table.

B Use the absolute value of the charge on each ion as the subscript for theother ion. Reduce the subscripts to the lowest numbers to write theempirical formula. Check to make sure the empirical formula is electricallyneutral.

Solution:

a. B Using the absolute values of the charges on the ions as thesubscripts gives Ga3As3:

Reducing the subscripts to the smallest whole numbers gives theempirical formula GaAs, which is electrically neutral [+3 + (−3) =0]. Alternatively, we could recognize that Ga3+ and As3− havecharges of equal magnitude but opposite signs. One Ga3+ ionbalances the charge on one As3− ion, and a 1:1 compound willhave no net charge. Because we write subscripts only if thenumber is greater than 1, the empirical formula is GaAs. GaAs is



gallium arsenide, which is widely used in the electronics industryin transistors and other devices.

b. B Because Eu3+ has a charge of +3 and O2− has a charge of −2, a 1:1compound would have a net charge of +1. We must therefore findmultiples of the charges that cancel. We cross charges, using theabsolute value of the charge on one ion as the subscript for theother ion:

The subscript for Eu3+ is 2 (from O2−), and the subscript for O2− is3 (from Eu3+), giving Eu2O3; the subscripts cannot be reducedfurther. The empirical formula contains a positive charge of 2(+3)= +6 and a negative charge of 3(−2) = −6, for a net charge of 0. Thecompound Eu2O3 is neutral. Europium oxide is responsible forthe red color in television and computer screens.

c. A Because the charges on the ions are not given, we must firstdetermine the charges expected for the most common ionsderived from calcium and chlorine. Calcium lies in group 2, so itshould lose two electrons to form Ca2+. Chlorine lies in group 17,so it should gain one electron to form Cl−.

B Two Cl− ions are needed to balance the charge on one Ca2+ ion,which leads to the empirical formula CaCl2. We could also crosscharges, using the absolute value of the charge on Ca2+ as thesubscript for Cl and the absolute value of the charge on Cl− as thesubscript for Ca:

The subscripts in CaCl2 cannot be reduced further. The empiricalformula is electrically neutral [+2 + 2(−1) = 0]. This compound iscalcium chloride, one of the substances used as “salt” to melt iceon roads and sidewalks in winter.

Exercise



Write the empirical formula for the simplest binary ionic compound formedfrom each ion or element pair.

a. Li+ and N3−

b. Al3+ and O2−

c. lithium and oxygen

Answer:

a. Li3Nb. Al2O3

c. Li2O

Polyatomic Ions

Polyatomic ions23 are groups of atoms that bear a net electrical charge, althoughthe atoms in a polyatomic ion are held together by the same covalent bonds thathold atoms together in molecules. Just as there are many more kinds of moleculesthan simple elements, there are many more kinds of polyatomic ions thanmonatomic ions. Two examples of polyatomic cations are the ammonium (NH4

+)

and the methylammonium (CH3NH3+) ions. Polyatomic anions are much more

numerous than polyatomic cations; some common examples are in Table 2.4"Common Polyatomic Ions and Their Names".

Table 2.4 Common Polyatomic Ions and Their Names

Formula Name of Ion

NH4+ ammonium

CH3NH3+ methylammonium

OH− hydroxide

O22− peroxide

CN− cyanide

SCN− thiocyanate

NO2− nitrite

NO3− nitrate

CO32− carbonate

23. A group of two or more atomsthat has a net electrical charge.



Formula Name of Ion

HCO3− hydrogen carbonate, or bicarbonate

SO32− sulfite

SO42− sulfate

HSO4− hydrogen sulfate, or bisulfate

PO43− phosphate

HPO42− hydrogen phosphate

H2PO4− dihydrogen phosphate

ClO− hypochlorite

ClO2− chlorite

ClO3− chlorate

ClO4− perchlorate

MnO4− permanganate

CrO42− chromate

Cr2O72− dichromate

C2O42− oxalate

HCO2− formate

CH3CO2− acetate

C6H5CO2− benzoate

The method we used to predict the empirical formulas for ionic compounds thatcontain monatomic ions can also be used for compounds that contain polyatomicions. The overall charge on the cations must balance the overall charge on theanions in the formula unit. Thus K+ and NO3

− ions combine in a 1:1 ratio to form

KNO3 (potassium nitrate or saltpeter), a major ingredient in black gunpowder.

Similarly, Ca2+ and SO42− form CaSO4 (calcium sulfate), which combines with

varying amounts of water to form gypsum and plaster of Paris. The polyatomic ionsNH4

+ and NO3− form NH4NO3 (ammonium nitrate), which is a widely used fertilizer

and, in the wrong hands, an explosive. One example of a compound in which theions have charges of different magnitudes is calcium phosphate, which is composedof Ca2+ and PO4

3− ions; it is a major component of bones. The compound is

electrically neutral because the ions combine in a ratio of three Ca2+ ions [3(+2) = +6]for every two ions [2(−3) = −6], giving an empirical formula of Ca3(PO4)2; the



parentheses around PO4 in the empirical formula indicate that it is a polyatomic

ion. Writing the formula for calcium phosphate as Ca3P2O8 gives the correct

number of each atom in the formula unit, but it obscures the fact that thecompound contains readily identifiable PO4

3− ions.



EXAMPLE 5

Write the empirical formula for the compound formed from each ion pair.

a. Na+ and HPO42−

b. potassium cation and cyanide anionc. calcium cation and hypochlorite anion

Given: ions

Asked for: empirical formula for ionic compound

Strategy:

A If it is not given, determine the charge on a monatomic ion from itslocation in the periodic table. Use Table 2.4 "Common Polyatomic Ions andTheir Names" to find the charge on a polyatomic ion.

B Use the absolute value of the charge on each ion as the subscript for theother ion. Reduce the subscripts to the smallest whole numbers whenwriting the empirical formula.

Solution:

a. B Because HPO42− has a charge of −2 and Na+ has a charge of +1, the

empirical formula requires two Na+ ions to balance the charge of thepolyatomic ion, giving Na2HPO4. The subscripts are reduced to thelowest numbers, so the empirical formula is Na2HPO4. This compound issodium hydrogen phosphate, which is used to provide texture inprocessed cheese, puddings, and instant breakfasts.

b. A The potassium cation is K+, and the cyanide anion is CN−. B Becausethe magnitude of the charge on each ion is the same, the empiricalformula is KCN. Potassium cyanide is highly toxic, and at one time it wasused as rat poison. This use has been discontinued, however, because toomany people were being poisoned accidentally.

c. A The calcium cation is Ca2+, and the hypochlorite anion is ClO−. B TwoClO− ions are needed to balance the charge on one Ca2+ ion, givingCa(ClO)2. The subscripts cannot be reduced further, so the empiricalformula is Ca(ClO)2. This is calcium hypochlorite, the “chlorine” used topurify water in swimming pools.



Exercise

Write the empirical formula for the compound formed from each ion pair.

a. Ca2+ and H2PO4−

b. sodium cation and bicarbonate anionc. ammonium cation and sulfate anion

Answer:

a. Ca(H2PO4)2: calcium dihydrogen phosphate is one of the ingredients inbaking powder.

b. NaHCO3: sodium bicarbonate is found in antacids and baking powder; inpure form, it is sold as baking soda.

c. (NH4)2SO4: ammonium sulfate is a common source of nitrogen infertilizers.

Hydrates

Many ionic compounds occur as hydrates24, compounds that contain specific ratiosof loosely bound water molecules, called waters of hydration25. Waters ofhydration can often be removed simply by heating. For example, calciumdihydrogen phosphate can form a solid that contains one molecule of water perCa(H2PO4)2 unit and is used as a leavening agent in the food industry to cause baked

goods to rise. The empirical formula for the solid is Ca(H2PO4)2·H2O. In contrast,

copper sulfate usually forms a blue solid that contains five waters of hydration performula unit, with the empirical formula CuSO4·5H2O. When heated, all five water

molecules are lost, giving a white solid with the empirical formula CuSO4 (Figure 2.9

"Loss of Water from a Hydrate with Heating").

24. A compound that containsspecific ratios of loosely boundwater molecules, called watersof hydration.

25. The loosely bound watermolecules in hydratecompounds. These waters ofhydration can often beremoved by simply heating thecompound.



Figure 2.9 Loss of Water from a Hydrate with Heating

When blue CuSO4·5H2O is heated, two molecules of water are lost at 30°C, two more at 110°C, and the last at 250°C to

give white CuSO4.

Compounds that differ only in the numbers of waters of hydration can have verydifferent properties. For example, CaSO4·½H2O is plaster of Paris, which was often

used to make sturdy casts for broken arms or legs, whereas CaSO4·2H2O is the less

dense, flakier gypsum, a mineral used in drywall panels for home construction.When a cast would set, a mixture of plaster of Paris and water crystallized to givesolid CaSO4·2H2O. Similar processes are used in the setting of cement and concrete.



Summary

An empirical formula gives the relative numbers of atoms of the elements in acompound, reduced to the lowest whole numbers. The formula unit is theabsolute grouping represented by the empirical formula of a compound, eitherionic or covalent. Empirical formulas are particularly useful for describing thecomposition of ionic compounds, which do not contain readily identifiablemolecules. Some ionic compounds occur as hydrates, which contain specificratios of loosely bound water molecules called waters of hydration.

KEY TAKEAWAY

• The composition of a compound is represented by an empirical ormolecular formula, each consisting of at least one formula unit.



CONCEPTUAL PROBLEMS

1. What are the differences and similarities between a polyatomic ion and amolecule?

2. Classify each compound as ionic or covalent.

a. Zn3(PO4)2b. C6H5CO2Hc. K2Cr2O7d. CH3CH2SHe. NH4Brf. CCl2F2

3. Classify each compound as ionic or covalent. Which are organic compoundsand which are inorganic compounds?

a. CH3CH2CO2Hb. CaCl2c. Y(NO3)3d. H2Se. NaC2H3O2

4. Generally, one cannot determine the molecular formula directly from anempirical formula. What other information is needed?

5. Give two pieces of information that we obtain from a structural formula thatwe cannot obtain from an empirical formula.

6. The formulas of alcohols are often written as ROH rather than as empiricalformulas. For example, methanol is generally written as CH3OH rather thanCH4O. Explain why the ROH notation is preferred.

7. The compound dimethyl sulfide has the empirical formula C2H6S and thestructural formula CH3SCH3. What information do we obtain from thestructural formula that we do not get from the empirical formula? Write thecondensed structural formula for the compound.

8. What is the correct formula for magnesium hydroxide—MgOH2 or Mg(OH)2?Why?

9. Magnesium cyanide is written as Mg(CN)2, not MgCN2. Why?

10. Does a given hydrate always contain the same number of waters of hydration?



ANSWER

7. The structural formula gives us the connectivity of the atoms in the moleculeor ion, as well as a schematic representation of their arrangement in space.Empirical formulas tell us only the ratios of the atoms present. The condensedstructural formula of dimethylsulfide is (CH3)2S.



NUMERICAL PROBLEMS

1. Write the formula for each compound.

a. magnesium sulfate, which has 1 magnesium atom, 4 oxygen atoms, and 1sulfur atom

b. ethylene glycol (antifreeze), which has 6 hydrogen atoms, 2 carbon atoms,and 2 oxygen atoms

c. acetic acid, which has 2 oxygen atoms, 2 carbon atoms, and 4 hydrogenatoms

d. potassium chlorate, which has 1 chlorine atom, 1 potassium atom, and 3oxygen atoms

e. sodium hypochlorite pentahydrate, which has 1 chlorine atom, 1 sodiumatom, 6 oxygen atoms, and 10 hydrogen atoms


a. cadmium acetate, which has 1 cadmium atom, 4 oxygen atoms, 4 carbonatoms, and 6 hydrogen atoms

b. barium cyanide, which has 1 barium atom, 2 carbon atoms, and 2 nitrogenatoms

c. iron(III) phosphate dihydrate, which has 1 iron atom, 1 phosphorus atom, 6oxygen atoms, and 4 hydrogen atoms

d. manganese(II) nitrate hexahydrate, which has 1 manganese atom, 12hydrogen atoms, 12 oxygen atoms, and 2 nitrogen atoms

e. silver phosphate, which has 1 phosphorus atom, 3 silver atoms, and 4oxygen atoms

3. Complete the following table by filling in the formula for the ionic compoundformed by each cation-anion pair.

Ion K+ Fe3+ NH4+ Ba2+

Cl− KCl

SO42−

PO43−

NO3−

OH−

4. Write the empirical formula for the binary compound formed by the mostcommon monatomic ions formed by each pair of elements.

a. zinc and sulfur



b. barium and iodinec. magnesium and chlorined. silicon and oxygene. sodium and sulfur

5. Write the empirical formula for the binary compound formed by the mostcommon monatomic ions formed by each pair of elements.

a. lithium and nitrogenb. cesium and chlorinec. germanium and oxygend. rubidium and sulfure. arsenic and sodium

6. Write the empirical formula for each compound.

a. Na2S2O4b. B2H6c. C6H12O6d. P4O10e. KMnO4

7. Write the empirical formula for each compound.

a. Al2Cl6b. K2Cr2O7c. C2H4d. (NH2)2CNHe. CH3COOH



ANSWERS

1. a. MgSO4b. C2H6O2c. C2H4O2d. KClO3e. NaOCl·5H2O

3. Ion K + Fe 3+ NH 4+ Ba 2+

Cl − KCl FeCl3 NH4Cl BaCl2

SO 42− K2SO4 Fe2(SO4)3 (NH4)2SO4 BaSO4

PO 43− K3PO4 FePO4 (NH4)3PO4 Ba3(PO4)2

NO 3− KNO3 Fe(NO3)3 NH4NO3 Ba(NO3)2

OH − KOH Fe(OH)3 NH4OH Ba(OH)2

5. a. Li3Nb. CsClc. GeO2d. Rb2Se. Na3As

7. a. AlCl3b. K2Cr2O7c. CH2d. CH5N3e. CH2O



2.3 Naming Ionic Compounds

LEARNING OBJECTIVE

1. To name ionic compounds.

The empirical and molecular formulas discussed in the preceding section areprecise and highly informative, but they have some disadvantages. First, they areinconvenient for routine verbal communication. For example, saying “C-A-three-P-O-four-two” for Ca3(PO4)2 is much more difficult than saying “calcium phosphate.”

In addition, you will see in Section 2.4 "Naming Covalent Compounds" that manycompounds have the same empirical and molecular formulas but differentarrangements of atoms, which result in very different chemical and physicalproperties. In such cases, it is necessary for the compounds to have different namesthat distinguish among the possible arrangements.

Many compounds, particularly those that have been known for a relatively longtime, have more than one name: a common name (sometimes more than one) and asystematic name, which is the name assigned by adhering to specific rules. Like thenames of most elements, the common names of chemical compounds generallyhave historical origins, although they often appear to be unrelated to thecompounds of interest. For example, the systematic name for KNO3 is potassium

nitrate, but its common name is saltpeter.

In this text, we use a systematic nomenclature to assign meaningful names to themillions of known substances. Unfortunately, some chemicals that are widely usedin commerce and industry are still known almost exclusively by their commonnames; in such cases, you must be familiar with the common name as well as thesystematic one. The objective of this and the next two sections is to teach you towrite the formula for a simple inorganic compound from its name—and viceversa—and introduce you to some of the more frequently encountered commonnames.

We begin with binary ionic compounds, which contain only two elements. Theprocedure for naming such compounds is outlined in Figure 2.10 "Naming an IonicCompound" and uses the following steps:


166

Figure 2.10 Naming an Ionic Compound

1. Place the ions in their proper order: cation and then anion.

2. Name the cation.

a. Metals that form only one cation. As noted in Section 2.1"Chemical Compounds", these metals are usually in groups 1–3, 12,and 13. The name of the cation of a metal that forms only onecation is the same as the name of the metal (with the word ionadded if the cation is by itself). For example, Na+ is the sodium ion,Ca2+ is the calcium ion, and Al3+ is the aluminum ion.

b. Metals that form more than one cation. As shown in Figure 2.11"Metals That Form More Than One Cation and Their Locations inthe Periodic Table", many metals can form more than one cation.This behavior is observed for most transition metals, manyactinides, and the heaviest elements of groups 13–15. In such cases,the positive charge on the metal is indicated by a roman numeralin parentheses immediately following the name of the metal. ThusCu+ is copper(I) (read as “copper one”), Fe2+ is iron(II), Fe3+ isiron(III), Sn2+ is tin(II), and Sn4+ is tin(IV).

An older system of nomenclature for such cations is still widelyused, however. The name of the cation with the higher charge isformed from the root of the element’s Latin name with the suffix-ic attached, and the name of the cation with the lower charge hasthe same root with the suffix -ous. The names of Fe3+, Fe2+, Sn4+,and Sn2+ are therefore ferric, ferrous, stannic, and stannous,respectively. Even though this text uses the systematic names withroman numerals, you should be able to recognize these commonnames because they are still often used. For example, on the label


2.3 Naming Ionic Compounds 167

of your dentist’s fluoride rinse, the compound chemists call tin(II)fluoride is usually listed as stannous fluoride.

Some examples of metals that form more than one cation are inTable 2.5 "Common Cations of Metals That Form More Than OneIon" along with the names of the ions. Note that the simple Hg+

cation does not occur in chemical compounds. Instead, allcompounds of mercury(I) contain a dimeric cation, Hg2

2+, in which

the two Hg atoms are bonded together.

Table 2.5 Common Cations of Metals That Form More Than One Ion

CationSystematic

NameCommon

NameCation

SystematicName

CommonName

Cr2+ chromium(II) chromous Cu2+ copper(II) cupric

Cr3+ chromium(III) chromic Cu+ copper(I) cuprous

Mn2+ manganese(II) manganous* Hg2+ mercury(II) mercuric

Mn3+ manganese(III) manganic* Hg22+ mercury(I) mercurous†

Fe2+ iron(II) ferrous Sn4+ tin(IV) stannic

Fe3+ iron(III) ferric Sn2+ tin(II) stannous

Co2+ cobalt(II) cobaltous* Pb4+ lead(IV) plumbic*

Co3+ cobalt(III) cobaltic* Pb2+ lead(II) plumbous*

* Not widely used.

†The isolated mercury(I) ion exists only as the gaseous ion.

c. Polyatomic cations. The names of the common polyatomic cationsthat are relatively important in ionic compounds (such as, theammonium ion) are in Table 2.4 "Common Polyatomic Ions andTheir Names".

3. Name the anion.

a. Monatomic anions. Monatomic anions are named by adding thesuffix -ide to the root of the name of the parent element; thus, Cl− ischloride, O2− is oxide, P3− is phosphide, N3− is nitride (also calledazide), and C4− is carbide. Because the charges on these ions can bepredicted from their position in the periodic table, it is notnecessary to specify the charge in the name. Examples of



monatomic anions are in Table 2.2 "Some Common MonatomicIons and Their Names".

b. Polyatomic anions. Polyatomic anions typically have commonnames that you must learn; some examples are in Table 2.4"Common Polyatomic Ions and Their Names". Polyatomic anionsthat contain a single metal or nonmetal atom plus one or moreoxygen atoms are called oxoanions (or oxyanions). In cases whereonly two oxoanions are known for an element, the name of theoxoanion with more oxygen atoms ends in -ate, and the name ofthe oxoanion with fewer oxygen atoms ends in -ite. For example,NO3− is nitrate and NO2− is nitrite.

The halogens and some of the transition metals form moreextensive series of oxoanions with as many as four members. In thenames of these oxoanions, the prefix per- is used to identify theoxoanion with the most oxygen (so that ClO4

− is perchlorate and

ClO3− is chlorate), and the prefix hypo- is used to identify the anion

with the fewest oxygen (ClO2− is chlorite and ClO− is hypochlorite).

The relationship between the names of oxoanions and the numberof oxygen atoms present is diagrammed in Figure 2.12 "TheRelationship between the Names of Oxoanions and the Number ofOxygen Atoms Present". Differentiating the oxoanions in such aseries is no trivial matter. For example, the hypochlorite ion is theactive ingredient in laundry bleach and swimming pooldisinfectant, but compounds that contain the perchlorate ion canexplode if they come into contact with organic substances.

4. Write the name of the compound as the name of the cation followed bythe name of the anion.

It is not necessary to indicate the number of cations or anions presentper formula unit in the name of an ionic compound because thisinformation is implied by the charges on the ions. You must considerthe charge of the ions when writing the formula for an ionic compoundfrom its name, however. Because the charge on the chloride ion is −1and the charge on the calcium ion is +2, for example, consistent withtheir positions in the periodic table, simple arithmetic tells you thatcalcium chloride must contain twice as many chloride ions as calciumions to maintain electrical neutrality. Thus the formula is CaCl2.

Similarly, calcium phosphate must be Ca3(PO4)2 because the cation and

the anion have charges of +2 and −3, respectively. The best way to learnhow to name ionic compounds is to work through a few examples,referring to Figure 2.10 "Naming an Ionic Compound", Table 2.2 "Some



Common Monatomic Ions and Their Names", Table 2.4 "CommonPolyatomic Ions and Their Names", and Table 2.5 "Common Cations ofMetals That Form More Than One Ion" as needed.

Figure 2.11 Metals That Form More Than One Cation and Their Locations in the Periodic Table

With only a few exceptions, these metals are usually transition metals or actinides.



Figure 2.12 The Relationship between the Names of Oxoanions and the Number of Oxygen Atoms Present

Note the Pattern

Cations are always named before anions.

Most transition metals, many actinides, and the heaviest elements of groups13–15 can form more than one cation.



EXAMPLE 6

Write the systematic name (and the common name if applicable) for eachionic compound.

a. LiClb. MgSO4

c. (NH4)3PO4

d. Cu2O

Given: empirical formula

Asked for: name

Strategy:

A If only one charge is possible for the cation, give its name, consultingTable 2.2 "Some Common Monatomic Ions and Their Names" or Table 2.4"Common Polyatomic Ions and Their Names" if necessary. If the cation canhave more than one charge (Table 2.5 "Common Cations of Metals That FormMore Than One Ion"), specify the charge using roman numerals.

B If the anion does not contain oxygen, name it according to step 3a, usingTable 2.2 "Some Common Monatomic Ions and Their Names" and Table 2.4"Common Polyatomic Ions and Their Names" if necessary. For polyatomicanions that contain oxygen, use Table 2.4 "Common Polyatomic Ions andTheir Names" and the appropriate prefix and suffix listed in step 3b.

C Beginning with the cation, write the name of the compound.

Solution:

a. A B Lithium is in group 1, so we know that it forms only the Li+ cation,which is the lithium ion. Similarly, chlorine is in group 7, so it forms theCl− anion, which is the chloride ion. C Because we begin with the nameof the cation, the name of this compound is lithium chloride, which isused medically as an antidepressant drug.

b. A B The cation is the magnesium ion, and the anion, which containsoxygen, is sulfate. C Because we list the cation first, the name of thiscompound is magnesium sulfate. A hydrated form of magnesium sulfate



(MgSO4·7H2O) is sold in drugstores as Epsom salts, a harsh but effectivelaxative.

c. A B The cation is the ammonium ion (from Table 2.4 "CommonPolyatomic Ions and Their Names"), and the anion is phosphate. C Thecompound is therefore ammonium phosphate, which is widely used as afertilizer. It is not necessary to specify that the formula unit containsthree ammonium ions because three are required to balance thenegative charge on phosphate.

d. A B The cation is a transition metal that often forms more than onecation (Table 2.5 "Common Cations of Metals That Form More Than OneIon"). We must therefore specify the positive charge on the cation in thename: copper(I) or, according to the older system, cuprous. The anion isoxide. C The name of this compound is copper(I) oxide or, in the oldersystem, cuprous oxide. Copper(I) oxide is used as a red glaze on ceramicsand in antifouling paints to prevent organisms from growing on thebottoms of boats.

Cu2O. The bottom of aboat is protected with ared antifouling paintcontaining copper(I)oxide, Cu2O.

Exercise

Write the systematic name (and the common name if applicable) for eachionic compound.

a. CuCl2



b. MgCO3

c. FePO4

Answer:

a. copper(II) chloride (or cupric chloride)b. magnesium carbonatec. iron(III) phosphate (or ferric phosphate)



EXAMPLE 7

Write the formula for each compound.

a. calcium dihydrogen phosphateb. aluminum sulfatec. chromium(III) oxide

Given: systematic name

Asked for: formula

Strategy:

A Identify the cation and its charge using the location of the element in theperiodic table and Table 2.2 "Some Common Monatomic Ions and TheirNames", Table 2.3 "The Physical Properties of Typical Ionic Compounds andCovalent Molecular Substances", Table 2.4 "Common Polyatomic Ions andTheir Names", and Table 2.5 "Common Cations of Metals That Form MoreThan One Ion". If the cation is derived from a metal that can form cationswith different charges, use the appropriate roman numeral or suffix toindicate its charge.

B Identify the anion using Table 2.2 "Some Common Monatomic Ions andTheir Names" and Table 2.4 "Common Polyatomic Ions and Their Names".Beginning with the cation, write the compound’s formula and thendetermine the number of cations and anions needed to achieve electricalneutrality.

Solution:

a. A Calcium is in group 2, so it forms only the Ca2+ ion. B Dihydrogenphosphate is the H2PO4

− ion (Table 2.4 "Common Polyatomic Ions andTheir Names"). Two H2PO4

− ions are needed to balance the positivecharge on Ca2+, to give Ca(H2PO4)2. A hydrate of calcium dihydrogenphosphate, Ca(H2PO4)2·H2O, is the active ingredient in baking powder.

b. A Aluminum, near the top of group 13 in the periodic table, forms onlyone cation, Al3+ (Figure 2.11 "Metals That Form More Than One Cationand Their Locations in the Periodic Table"). B Sulfate is SO4

2− (Table 2.4"Common Polyatomic Ions and Their Names"). To balance the electricalcharges, we need two Al3+ cations and three SO4

2− anions, giving



Al2(SO4)3. Aluminum sulfate is used to tan leather and purify drinkingwater.

c. A Because chromium is a transition metal, it can form cations withdifferent charges. The roman numeral tells us that the positive charge inthis case is +3, so the cation is Cr3+. B Oxide is O2−. Thus two cations(Cr3+) and three anions (O2−) are required to give an electrically neutralcompound, Cr2O3. This compound is a common green pigment that hasmany uses, including camouflage coatings.

Cr2O3. Chromium(III)oxide (Cr2O3) is acommon pigment indark green paints, suchas camouflage paint.

Exercise

Write the formula for each compound.

a. barium chlorideb. sodium carbonatec. iron(III) hydroxide

Answer:

a. BaCl2b. Na2CO3

c. Fe(OH)3



Summary

Ionic compounds are named according to systematic procedures, althoughcommon names are widely used. Systematic nomenclature enables us to writethe structure of any compound from its name and vice versa. Ionic compoundsare named by writing the cation first, followed by the anion. If a metal can formcations with more than one charge, the charge is indicated by roman numeralsin parentheses following the name of the metal. Oxoanions are polyatomicanions that contain a single metal or nonmetal atom and one or more oxygenatoms.

KEY TAKEAWAY

• There is a systematic method used to name ionic compounds.



CONCEPTUAL PROBLEMS

1. Name each cation.

a. K+

b. Al3+

c. NH4+

d. Mg2+

e. Li+

2. Name each anion.

a. Br−

b. CO32−

c. S2−

d. NO3−

e. HCO2−

f. F−

g. ClO−

h. C2O42−

3. Name each anion.

a. PO43−

b. Cl−

c. SO32−

d. CH3CO2−

e. HSO4−

f. ClO4−

g. NO2−

h. O2−

4. Name each anion.

a. SO42−

b. CN−

c. Cr2O72−

d. N3−

e. OH−

f. I−

g. O22−

5. Name each compound.



a. MgBr2b. NH4CNc. CaOd. KClO3e. K3PO4f. NH4NO2g. NaN3


a. NaNO3b. Cu3(PO4)2c. NaOHd. Li4Ce. CaF2f. NH4Brg. MgCO3


a. RbBrb. Mn2(SO4)3c. NaClOd. (NH4)2SO4e. NaBrf. KIO3g. Na2CrO4


a. NH4ClO4b. SnCl4c. Fe(OH)2d. Na2Oe. MgCl2f. K2SO4g. RaCl2


a. KCNb. LiOHc. CaCl2



d. NiSO4e. NH4ClO2f. LiClO4g. La(CN)3

ANSWER

7. a. rubidium bromideb. manganese(III) sulfatec. sodium hypochlorited. ammonium sulfatee. sodium bromidef. potassium iodateg. sodium chromate



NUMERICAL PROBLEMS

1. For each ionic compound, name the cation and the anion and give the chargeon each ion.

a. BeOb. Pb(OH)2c. BaSd. Na2Cr2O7e. ZnSO4f. KClOg. NaH2PO4

2. For each ionic compound, name the cation and the anion and give the chargeon each ion.

a. Zn(NO3)2b. CoSc. BeCO3d. Na2SO4e. K2C2O4f. NaCNg. FeCl2


a. magnesium carbonateb. aluminum sulfatec. potassium phosphated. lead(IV) oxidee. silicon nitridef. sodium hypochloriteg. titanium(IV) chlorideh. disodium ammonium phosphate


a. lead(II) nitrateb. ammonium phosphatec. silver sulfided. barium sulfatee. cesium iodidef. sodium bicarbonateg. potassium dichromate



h. sodium hypochlorite


a. zinc cyanideb. silver chromatec. lead(II) iodided. benzenee. copper(II) perchlorate


a. calcium fluorideb. sodium nitratec. iron(III) oxided. copper(II) acetatee. sodium nitrite


a. sodium hydroxideb. calcium cyanidec. magnesium phosphated. sodium sulfatee. nickel(II) bromidef. calcium chloriteg. titanium(IV) bromide


a. sodium chloriteb. potassium nitritec. sodium nitride (also called sodium azide)d. calcium phosphidee. tin(II) chloridef. calcium hydrogen phosphateg. iron(II) chloride dihydrate


a. potassium carbonateb. chromium(III) sulfitec. cobalt(II) phosphated. magnesium hypochloritee. nickel(II) nitrate hexahydrate



Figure 2.13 Naming aCovalent InorganicCompound

2.4 Naming Covalent Compounds

LEARNING OBJECTIVE

1. To name covalent compounds that contain up to three elements.

As with ionic compounds, the system that chemists have devised for namingcovalent compounds enables us to write the molecular formula from the name andvice versa. In this and the following section, we describe the rules for namingsimple covalent compounds. We begin with inorganic compounds and then turn tosimple organic compounds that contain only carbon and hydrogen.

Binary Inorganic Compounds

Binary covalent compounds—that is, covalent compounds that contain only twoelements—are named using a procedure similar to that used to name simple ioniccompounds, but prefixes are added as needed to indicate the number of atoms ofeach kind. The procedure, diagrammed in Figure 2.13 "Naming a Covalent InorganicCompound", uses the following steps:

1. Place the elements in their proper order.

a. The element farthest to the left in theperiodic table is usually named first. Ifboth elements are in the same group,the element closer to the bottom of thecolumn is named first.

b. The second element is named as if itwere a monatomic anion in an ioniccompound (even though it is not), withthe suffix -ide attached to the root ofthe element name.

2. Identify the number of each type of atompresent.

a. Prefixes derived from Greek stems are used to indicate the numberof each type of atom in the formula unit (Table 2.6 "Prefixes forIndicating the Number of Atoms in Chemical Names"). The prefix


183

mono- (“one”) is used only when absolutely necessary to avoidconfusion, just as we omit the subscript 1 when writing molecularformulas.

To demonstrate steps 1 and 2a, we name HCl as hydrogen chloride(because hydrogen is to the left of chlorine in the periodic table)and PCl5 as phosphorus pentachloride. The order of the elements

in the name of BrF3, bromine trifluoride, is determined by the fact

that bromine lies below fluorine in group 17.

Table 2.6 Prefixes for Indicating the Number of Atoms in ChemicalNames

Prefix Number

mono- 1

di- 2

tri- 3

tetra- 4

penta- 5

hexa- 6

hepta- 7

octa- 8

nona- 9

deca- 10

undeca- 11

dodeca- 12

b. If a molecule contains more than one atom of both elements, thenprefixes are used for both. Thus N2O3 is dinitrogen trioxide, as

shown in Figure 2.13 "Naming a Covalent Inorganic Compound".c. In some names, the final a or o of the prefix is dropped to avoid

awkward pronunciation. Thus OsO4 is osmium tetroxide rather

than osmium tetraoxide.

3. Write the name of the compound.


2.4 Naming Covalent Compounds 184

a. Binary compounds of the elements with oxygen are generallynamed as “element oxide,” with prefixes that indicate the numberof atoms of each element per formula unit. For example, CO iscarbon monoxide. The only exception is binary compounds ofoxygen with fluorine, which are named as oxygen fluorides. (Thereasons for this convention will become clear in Chapter 7 "ThePeriodic Table and Periodic Trends" and Chapter 8 "Ionic versusCovalent Bonding".)

b. Certain compounds are always called by the common names thatwere assigned long ago when names rather than formulas wereused. For example, H2O is water (not dihydrogen oxide); NH3 is

ammonia; PH3 is phosphine; SiH4 is silane; and B2H6, a dimer of

BH3, is diborane. For many compounds, the systematic name and

the common name are both used frequently, so you must befamiliar with them. For example, the systematic name for NO isnitrogen monoxide, but it is much more commonly called nitricoxide. Similarly, N2O is usually called nitrous oxide rather than

dinitrogen monoxide. Notice that the suffixes -ic and -ous are thesame ones used for ionic compounds.

Note the Pattern

Start with the element at the far left in the periodic table and work to the right.If two or more elements are in the same group, start with the bottom elementand work up.



EXAMPLE 8

Write the name of each binary covalent compound.

a. SF6

b. N2O4

c. ClO2

Given: molecular formula

Asked for: name of compound

Strategy:

A List the elements in order according to their positions in the periodictable. Identify the number of each type of atom in the chemical formula andthen use Table 2.6 "Prefixes for Indicating the Number of Atoms in ChemicalNames" to determine the prefixes needed.

B If the compound contains oxygen, follow step 3a. If not, decide whether touse the common name or the systematic name.

Solution:

a. A Because sulfur is to the left of fluorine in the periodic table, sulfur isnamed first. Because there is only one sulfur atom in the formula, noprefix is needed. B There are, however, six fluorine atoms, so we use theprefix for six: hexa- (Table 2.6 "Prefixes for Indicating the Number ofAtoms in Chemical Names"). The compound is sulfur hexafluoride.

b. A Because nitrogen is to the left of oxygen in the periodic table, nitrogenis named first. Because more than one atom of each element is present,prefixes are needed to indicate the number of atoms of each. Accordingto Table 2.6 "Prefixes for Indicating the Number of Atoms in ChemicalNames", the prefix for two is di-, and the prefix for four is tetra-. B Thecompound is dinitrogen tetroxide (omitting the a in tetra- according tostep 2c) and is used as a component of some rocket fuels.

c. A Although oxygen lies to the left of chlorine in the periodic table, it isnot named first because ClO2 is an oxide of an element other thanfluorine (step 3a). Consequently, chlorine is named first, but a prefix isnot necessary because each molecule has only one atom of chlorine. BBecause there are two oxygen atoms, the compound is a dioxide. Thus



the compound is chlorine dioxide. It is widely used as a substitute forchlorine in municipal water treatment plants because, unlike chlorine, itdoes not react with organic compounds in water to produce potentiallytoxic chlorinated compounds.

Exercise

Write the name of each binary covalent compound.

a. IF7

b. N2O5

c. OF2

Answer:

a. iodine heptafluorideb. dinitrogen pentoxidec. oxygen difluoride



EXAMPLE 9

Write the formula for each binary covalent compound.

a. sulfur trioxideb. diiodine pentoxide

Given: name of compound

Asked for: formula

Strategy:

List the elements in the same order as in the formula, use Table 2.6 "Prefixesfor Indicating the Number of Atoms in Chemical Names" to identify thenumber of each type of atom present, and then indicate this quantity as asubscript to the right of that element when writing the formula.

Solution:

a. Sulfur has no prefix, which means that each molecule has only onesulfur atom. The prefix tri- indicates that there are three oxygen atoms.The formula is therefore SO3. Sulfur trioxide is produced industrially inhuge amounts as an intermediate in the synthesis of sulfuric acid.

b. The prefix di- tells you that each molecule has two iodine atoms, and theprefix penta- indicates that there are five oxygen atoms. The formula isthus I2O5, a compound used to remove carbon monoxide from air inrespirators.

Exercise

Write the formula for each binary covalent compound.

a. silicon tetrachlorideb. disulfur decafluoride

Answer:

a. SiCl4b. S2F10



The structures of some of the compounds in Example 8 and Example 9 are shown inFigure 2.14 "The Structures of Some Covalent Inorganic Compounds and theLocations of the “Central Atoms” in the Periodic Table", along with the location ofthe “central atom” of each compound in the periodic table. It may seem that thecompositions and structures of such compounds are entirely random, but this is nottrue. After you have mastered the material in Chapter 7 "The Periodic Table andPeriodic Trends" and Chapter 8 "Ionic versus Covalent Bonding", you will be able topredict the compositions and structures of compounds of this type with a highdegree of accuracy.

Figure 2.14 The Structures of Some Covalent Inorganic Compounds and the Locations of the “Central Atoms” inthe Periodic Table

The compositions and structures of covalent inorganic compounds are not random. As you will learn in Chapter 7"The Periodic Table and Periodic Trends" and Chapter 8 "Ionic versus Covalent Bonding", they can be predictedfrom the locations of the component atoms in the periodic table.

Hydrocarbons

Approximately one-third of the compounds produced industrially are organiccompounds. All living organisms are composed of organic compounds, as is most ofthe food you consume, the medicines you take, the fibers in the clothes you wear,and the plastics in the materials you use. Section 2.1 "Chemical Compounds"introduced two organic compounds: methane (CH4) and methanol (CH3OH). These



and other organic compounds appear frequently in discussions and examplesthroughout this text.

The detection of organic compounds is useful in many fields. In one recentlydeveloped application, scientists have devised a new method called “materialdegradomics” to make it possible to monitor the degradation of old books andhistorical documents. As paper ages, it produces a familiar “old book smell” fromthe release of organic compounds in gaseous form. The composition of the gasdepends on the original type of paper used, a book’s binding, and the applied media.By analyzing these organic gases and isolating the individual components,preservationists are better able to determine the condition of an object and thosebooks and documents most in need of immediate protection.

The simplest class of organic compounds is the hydrocarbons26, which consistentirely of carbon and hydrogen. Petroleum and natural gas are complex, naturallyoccurring mixtures of many different hydrocarbons that furnish raw materials forthe chemical industry. The four major classes of hydrocarbons are the alkanes27,which contain only carbon–hydrogen and carbon–carbon single bonds; thealkenes28, which contain at least one carbon–carbon double bond; the alkynes29,which contain at least one carbon–carbon triple bond; and the aromatichydrocarbons30, which usually contain rings of six carbon atoms that can be drawnwith alternating single and double bonds. Alkanes are also called saturatedhydrocarbons, whereas hydrocarbons that contain multiple bonds (alkenes,alkynes, and aromatics) are unsaturated.

Alkanes

The simplest alkane is methane (CH4), a colorless, odorless gas that is the major

component of natural gas. In larger alkanes whose carbon atoms are joined in anunbranched chain (straight-chain alkanes), each carbon atom is bonded to at mosttwo other carbon atoms. The structures of two simple alkanes are shown in Figure2.15 "Straight-Chain Alkanes with Two and Three Carbon Atoms", and the namesand condensed structural formulas for the first 10 straight-chain alkanes are inTable 2.7 "The First 10 Straight-Chain Alkanes". The names of all alkanes end in-ane, and their boiling points increase as the number of carbon atoms increases.

26. The simplest class of organicmolecules, consisting of onlycarbon and hydrogen.

27. A saturated hydrocarbon withonly carbon–hydrogen andcarbon–carbon single bonds.

28. An unsaturated hydrocarbonwith at least onecarbon–carbon double bond.

29. An unsaturated hydrocarbonwith at least onecarbon–carbon triple bond.

30. An unsaturated hydrocarbonconsisting of a ring of sixcarbon atoms with alternatingsingle and double bonds.



Figure 2.15 Straight-Chain Alkanes with Two and Three Carbon Atoms

Table 2.7 The First 10 Straight-Chain Alkanes

NameNumber of

Carbon AtomsMolecularFormula

CondensedStructuralFormula

BoilingPoint (°C)

Uses

methane 1 CH4 CH4 −162natural gasconstituent

ethane 2 C2H6 CH3CH3 −89natural gasconstituent

propane 3 C3H8 CH3CH2CH3 −42 bottled gas

butane 4 C4H10CH3CH2CH2CH3 orCH3(CH2)2CH3

0lighters,bottled gas

pentane 5 C5H12 CH3(CH2)3CH3 36solvent,gasoline

hexane 6 C6H14 CH3(CH2)4CH3 69solvent,gasoline

heptane 7 C7H16 CH3(CH2)5CH3 98solvent,gasoline

octane 8 C8H18 CH3(CH2)6CH3 126 gasoline

nonane 9 C9H20 CH3(CH2)7CH3 151 gasoline

decane 10 C10H22 CH3(CH2)8CH3 174 kerosene



Alkanes with four or more carbon atoms can have more than one arrangement ofatoms. The carbon atoms can form a single unbranched chain, or the primary chainof carbon atoms can have one or more shorter chains that form branches. Forexample, butane (C4H10) has two possible structures. Normal butane (usually called

n-butane) is CH3CH2CH2CH3, in which the carbon atoms form a single unbranched

chain. In contrast, the condensed structural formula for isobutane is (CH3)2CHCH3, in

which the primary chain of three carbon atoms has a one-carbon chain branchingat the central carbon. Three-dimensional representations of both structures are asfollows:

The systematic names for branched hydrocarbons use the lowest possible numberto indicate the position of the branch along the longest straight carbon chain in thestructure. Thus the systematic name for isobutane is 2-methylpropane, whichindicates that a methyl group (a branch consisting of –CH3) is attached to the

second carbon of a propane molecule. Similarly, you will learn in Section 2.6"Industrially Important Chemicals" that one of the major components of gasoline iscommonly called isooctane; its structure is as follows:



As you can see, the compound has a chain of five carbon atoms, so it is a derivativeof pentane. There are two methyl group branches at one carbon atom and onemethyl group at another. Using the lowest possible numbers for the branches gives2,2,4-trimethylpentane for the systematic name of this compound.

Alkenes

The simplest alkenes are ethylene, C2H4 or CH2=CH2, and propylene, C3H6 or

CH3CH=CH2 (part (a) in Figure 2.16 "Some Simple (a) Alkenes, (b) Alkynes, and (c)

Cyclic Hydrocarbons"). The names of alkenes that have more than three carbonatoms use the same stems as the names of the alkanes (Table 2.7 "The First 10Straight-Chain Alkanes") but end in -ene instead of -ane.

Once again, more than one structure is possible for alkenes with four or morecarbon atoms. For example, an alkene with four carbon atoms has three possiblestructures. One is CH2=CHCH2CH3 (1-butene), which has the double bond between

the first and second carbon atoms in the chain. The other two structures have thedouble bond between the second and third carbon atoms and are forms ofCH3CH=CHCH3 (2-butene). All four carbon atoms in 2-butene lie in the same plane,

so there are two possible structures (part (a) in Figure 2.16 "Some Simple (a)



Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons"). If the two methyl groups are onthe same side of the double bond, the compound is cis-2-butene (from the Latin cis,meaning “on the same side”). If the two methyl groups are on opposite sides of thedouble bond, the compound is trans-2-butene (from the Latin trans, meaning“across”). These are distinctly different molecules: cis-2-butene melts at −138.9°C,whereas trans-2-butene melts at −105.5°C.

Figure 2.16 Some Simple (a) Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons

The positions of the carbon atoms in the chain are indicated by C1 or C2.

Just as a number indicates the positions of branches in an alkane, the number in thename of an alkene specifies the position of the first carbon atom of the double bond.The name is based on the lowest possible number starting from either end of thecarbon chain, so CH3CH2CH=CH2 is called 1-butene, not 3-butene. Note that

CH2=CHCH2CH3 and CH3CH2CH=CH2 are different ways of writing the same molecule

(1-butene) in two different orientations.



The name of a compound does not depend on its orientation. As illustrated for 1-butene, both condensedstructural formulas and molecular models show different orientations of the same molecule. Don’t let orientationfool you; you must be able to recognize the same structure no matter what its orientation.

Note the Pattern

The positions of groups or multiple bonds are always indicated by the lowestnumber possible.

Alkynes

The simplest alkyne is acetylene, C2H2 or HC≡CH (part (b) in Figure 2.16 "Some

Simple (a) Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons"). Because a mixture ofacetylene and oxygen burns with a flame that is hot enough (>3000°C) to cut metalssuch as hardened steel, acetylene is widely used in cutting and welding torches. Thenames of other alkynes are similar to those of the corresponding alkanes but end in



-yne. For example, HC≡CCH3 is propyne, and CH3C≡CCH3 is 2-butyne because the

multiple bond begins on the second carbon atom.

Note the Pattern

The number of bonds between carbon atoms in a hydrocarbon is indicated inthe suffix:

• alkane: only carbon–carbon single bonds• alkene: at least one carbon–carbon double bond• alkyne: at least one carbon–carbon triple bond

Cyclic Hydrocarbons

In a cyclic hydrocarbon31, the ends of a hydrocarbon chain are connected to form aring of covalently bonded carbon atoms. Cyclic hydrocarbons are named byattaching the prefix cyclo- to the name of the alkane, the alkene, or the alkyne. Thesimplest cyclic alkanes are cyclopropane (C3H6) a flammable gas that is also a

powerful anesthetic, and cyclobutane (C4H8) (part (c) in Figure 2.16 "Some Simple (a)

Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons"). The most common way to drawthe structures of cyclic alkanes is to sketch a polygon with the same number ofvertices as there are carbon atoms in the ring; each vertex represents a CH2 unit.

The structures of the cycloalkanes that contain three to six carbon atoms are shownschematically in Figure 2.17 "The Simple Cycloalkanes".

31. A hydrocarbon in which theends of the carbon chain areconnected to form a ring ofcovalently bonded carbonatoms.



Figure 2.17 The Simple Cycloalkanes

Aromatic Hydrocarbons

Alkanes, alkenes, alkynes, and cyclic hydrocarbons are generally called aliphatichydrocarbons32. The name comes from the Greek aleiphar, meaning “oil,” becausethe first examples were extracted from animal fats. In contrast, the first examplesof aromatic hydrocarbons, also called arenes, were obtained by the distillation anddegradation of highly scented (thus aromatic) resins from tropical trees.

The simplest aromatic hydrocarbon is benzene (C6H6), which was first obtained from

a coal distillate. The word aromatic now refers to benzene and structurally similarcompounds. As shown in part (a) in Figure 2.18 "Two Aromatic Hydrocarbons: (a)Benzene and (b) Toluene", it is possible to draw the structure of benzene in twodifferent but equivalent ways, depending on which carbon atoms are connected bydouble bonds or single bonds. Toluene is similar to benzene, except that onehydrogen atom is replaced by a –CH3 group; it has the formula C7H8 (part (b) in

Figure 2.18 "Two Aromatic Hydrocarbons: (a) Benzene and (b) Toluene"). As you willsoon learn, the chemical behavior of aromatic compounds differs from the behaviorof aliphatic compounds. Benzene and toluene are found in gasoline, and benzene isthe starting material for preparing substances as diverse as aspirin and nylon.

32. Alkanes, alkenes, alkynes, andcyclic hydrocarbons(hydrocarbons that are notaromatic).



Figure 2.18 Two Aromatic Hydrocarbons: (a) Benzene and (b) Toluene

Figure 2.19 "Two Hydrocarbons with the Molecular Formula C" illustrates two of themolecular structures possible for hydrocarbons that have six carbon atoms. As youcan see, compounds with the same molecular formula can have very differentstructures.

Figure 2.19 Two Hydrocarbons with the Molecular Formula C6H12



EXAMPLE 10

Write the condensed structural formula for each hydrocarbon.

a. n-heptaneb. 2-pentenec. 2-butyned. cyclooctene

Given: name of hydrocarbon

Asked for: condensed structural formula

Strategy:

A Use the prefix to determine the number of carbon atoms in the moleculeand whether it is cyclic. From the suffix, determine whether multiple bondsare present.

B Identify the position of any multiple bonds from the number(s) in thename and then write the condensed structural formula.

Solution:

a. A The prefix hept- tells us that this hydrocarbon has seven carbonatoms, and n- indicates that the carbon atoms form a straight chain. Thesuffix -ane tells that it is an alkane, with no carbon–carbon double ortriple bonds. B The condensed structural formula isCH3CH2CH2CH2CH2CH2CH3, which can also be written as CH3(CH2)5CH3.

b. A The prefix pent- tells us that this hydrocarbon has five carbonatoms, and the suffix -ene indicates that it is an alkene, with acarbon–carbon double bond. B The 2- tells us that the doublebond begins on the second carbon of the five-carbon atom chain.The condensed structural formula of the compound is thereforeCH3CH=CHCH2CH3.



c. A The prefix but- tells us that the compound has a chain of fourcarbon atoms, and the suffix -yne indicates that it has acarbon–carbon triple bond. B The 2- tells us that the triple bondbegins on the second carbon of the four-carbon atom chain. Sothe condensed structural formula for the compound isCH3C≡CCH3.



d. A The prefix cyclo- tells us that this hydrocarbon has a ringstructure, and oct- indicates that it contains eight carbon atoms,which we can draw as

The suffix -ene tells us that the compound contains acarbon–carbon double bond, but where in the ring do we placethe double bond? B Because all eight carbon atoms are identical,it doesn’t matter. We can draw the structure of cyclooctene as

Exercise

Write the condensed structural formula for each hydrocarbon.

a. n-octaneb. 2-hexenec. 1-heptyned. cyclopentane

Answer:

a. CH3(CH2)6CH3

b. CH3CH=CHCH2CH2CH3

c. HC≡C(CH2)4CH3

d.

The general name for a group of atoms derived from an alkane is an alkyl group. Thename of an alkyl group is derived from the name of the alkane by adding the suffix-yl. Thus the –CH3 fragment is a methyl group, the –CH2CH3 fragment is an ethyl



group, and so forth, where the dash represents a single bond to some other atom orgroup. Similarly, groups of atoms derived from aromatic hydrocarbons are arylgroups, which sometimes have unexpected names. For example, the –C6H5 fragment

is derived from benzene, but it is called a phenyl group. In general formulas andstructures, alkyl and aryl groups are often abbreviated as R33.

Structures of alkyl and aryl groups. The methyl group is an example of an alkyl group, and the phenyl group isan example of an aryl group.

Alcohols

Replacing one or more hydrogen atoms of a hydrocarbon with an –OH group givesan alcohol34, represented as ROH. The simplest alcohol (CH3OH) is called either

methanol (its systematic name) or methyl alcohol (its common name) (see Figure 2.4"Different Ways of Representing the Structure of a Molecule"). Methanol is theantifreeze in automobile windshield washer fluids, and it is also used as an efficientfuel for racing cars, most notably in the Indianapolis 500. Ethanol (or ethyl alcohol,CH3CH2OH) is familiar as the alcohol in fermented or distilled beverages, such as

beer, wine, and whiskey; it is also used as a gasoline additive (Section 2.6"Industrially Important Chemicals"). The simplest alcohol derived from an aromatic

33. The abbreviation used for alkylgroups and aryl groups ingeneral formulas andstructures.

34. A class of organic compoundsobtained by replacing one ormore of the hydrogen atoms ofa hydrocarbon with an −OHgroup.



hydrocarbon is C6H5OH, phenol (shortened from phenyl alcohol), a potent

disinfectant used in some sore throat medications and mouthwashes.

Ethanol, which is easy to obtain from fermentationprocesses, has successfully been used as an alternativefuel for several decades. Although it is a “green” fuelwhen derived from plants, it is an imperfect substitutefor fossil fuels because it is less efficient than gasoline.Moreover, because ethanol absorbs water from theatmosphere, it can corrode an engine’s seals. Thus othertypes of processes are being developed that use bacteriato create more complex alcohols, such as octanol, thatare more energy efficient and that have a lowertendency to absorb water. As scientists attempt toreduce mankind’s dependence on fossil fuels, thedevelopment of these so-called biofuels is a particularlyactive area of research.



Summary

Covalent inorganic compounds are named by a procedure similar to that usedfor ionic compounds, using prefixes to indicate the numbers of atoms in themolecular formula. The simplest organic compounds are the hydrocarbons,which contain only carbon and hydrogen. Alkanes contain onlycarbon–hydrogen and carbon–carbon single bonds, alkenes contain at least onecarbon–carbon double bond, and alkynes contain one or more carbon–carbontriple bonds. Hydrocarbons can also be cyclic, with the ends of the chainconnected to form a ring. Collectively, alkanes, alkenes, and alkynes are calledaliphatic hydrocarbons. Aromatic hydrocarbons, or arenes, are anotherimportant class of hydrocarbons that contain rings of carbon atoms related tothe structure of benzene (C6H6). A derivative of an alkane or an arene from

which one hydrogen atom has been removed is called an alkyl group or an arylgroup, respectively. Alcohols are another common class of organic compound,which contain an –OH group covalently bonded to either an alkyl group or anaryl group (often abbreviated R).

KEY TAKEAWAY

• Covalent inorganic compounds are named using a procedure similar tothat used for ionic compounds, whereas hydrocarbons use a systembased on the number of bonds between carbon atoms.



CONCEPTUAL PROBLEMS

1. Benzene (C6H6) is an organic compound, and KCl is an ionic compound. Thesum of the masses of the atoms in each empirical formula is approximately thesame. How would you expect the two to compare with regard to each of thefollowing? What species are present in benzene vapor?

a. melting pointb. type of bondingc. rate of evaporationd. structure

2. Can an inorganic compound be classified as a hydrocarbon? Why or why not?

3. Is the compound NaHCO3 a hydrocarbon? Why or why not?


a. NiOb. TiO2c. N2Od. CS2e. SO3f. NF3g. SF6


a. HgCl2b. IF5c. N2O5d. Cl2Oe. HgSf. PCl5

6. For each structural formula, write the condensed formula and the name of thecompound.

a.



b.

c.

d.

e.


a.

b.



c.

d.

8. Would you expect PCl3 to be an ionic compound or a covalent compound?Explain your reasoning.

9. What distinguishes an aromatic hydrocarbon from an aliphatic hydrocarbon?

10. The following general formulas represent specific classes of hydrocarbons.Refer to Table 2.7 "The First 10 Straight-Chain Alkanes" and Table 2.8 "SomeCommon Acids That Do Not Contain Oxygen" and Figure 2.16 "Some Simple (a)Alkenes, (b) Alkynes, and (c) Cyclic Hydrocarbons" and identify the classes.

a. CnH2n + 2b. CnH2nc. CnH2n − 2

11. Using R to represent an alkyl or aryl group, show the general structure of an

a. alcohol.b. phenol.

ANSWER

11. a. ROH (where R is an alkyl group)b. ROH (where R is an aryl group)



NUMERICAL PROBLEMS


a. dinitrogen monoxideb. silicon tetrafluoridec. boron trichlorided. nitrogen trifluoridee. phosphorus tribromide


a. dinitrogen trioxideb. iodine pentafluoridec. boron tribromided. oxygen difluoridee. arsenic trichloride


a. thallium(I) selenideb. neptunium(IV) oxidec. iron(II) sulfided. copper(I) cyanidee. nitrogen trichloride


a. RuO4b. PbO2c. MoF6d. Hg2(NO3)2·2H2Oe. WCl4


a. NbO2b. MoS2c. P4S10d. Cu2Oe. ReF5

6. Draw the structure of each compound.

a. propyneb. ethanol



c. n-hexaned. cyclopropanee. benzene

7. Draw the structure of each compound.

a. 1-buteneb. 2-pentynec. cycloheptaned. toluenee. phenol



ANSWERS

1. a. N2Ob. SiF4c. BCl3d. NF3e. PBr3

3. a. Tl2Seb. NpO2c. FeSd. CuCNe. NCl3

5. a. niobium (IV) oxideb. molybdenum (IV) sulfidec. tetraphosphorus decasulfided. copper(I) oxidee. rhenium(V) fluoride

7. a.

b.

c.

d.

e.



2.5 Acids and Bases

LEARNING OBJECTIVE

1. To identify and name some common acids and bases.

For our purposes at this point in the text, we can define an acid35 as a substancewith at least one hydrogen atom that can dissociate to form an anion and an H+ ion(a proton) in aqueous solution, thereby forming an acidic solution. We can definebases36 as compounds that produce hydroxide ions (OH−) and a cation whendissolved in water, thus forming a basic solution. Solutions that are neither basic noracidic are neutral. We will discuss the chemistry of acids and bases in more detail inChapter 4 "Reactions in Aqueous Solution", Chapter 8 "Ionic versus CovalentBonding", and Chapter 16 "Aqueous Acid–Base Equilibriums", but in this section wedescribe the nomenclature of common acids and identify some important bases sothat you can recognize them in future discussions. Pure acids and bases and theirconcentrated aqueous solutions are commonly encountered in the laboratory. Theyare usually highly corrosive, so they must be handled with care.

Acids

The names of acids differentiate between (1) acids in which the H+ ion is attached toan oxygen atom of a polyatomic anion (these are called oxoacids37, or occasionallyoxyacids) and (2) acids in which the H+ ion is attached to some other element. Inthe latter case, the name of the acid begins with hydro- and ends in -ic, with the rootof the name of the other element or ion in between. Recall that the name of theanion derived from this kind of acid always ends in -ide. Thus hydrogen chloride(HCl) gas dissolves in water to form hydrochloric acid (which contains H+ and Cl−

ions), hydrogen cyanide (HCN) gas forms hydrocyanic acid (which contains H+ andCN− ions), and so forth (Table 2.8 "Some Common Acids That Do Not ContainOxygen"). Examples of this kind of acid are commonly encountered and veryimportant. For instance, your stomach contains a dilute solution of hydrochloricacid to help digest food. When the mechanisms that prevent the stomach fromdigesting itself malfunction, the acid destroys the lining of the stomach and anulcer forms.

35. A substance with at least onehydrogen atom that candissociate to form an anion andan H+ ion (a proton) inaqueous solution, therebyfoming an acidic solution.

36. A substance that produces oneor more hydroxide ions(OH−) and a cation whendissolved in aqueous solution,thereby forming a basicsolution.

37. An acid in which thedissociable H+ ion is attachedto an oxygen atom of apolyatomic anion.


211

Note the Pattern

Acids are distinguished by whether the H+ ion is attached to an oxygen atom ofa polyatomic anion or some other element.

Table 2.8 Some Common Acids That Do Not Contain Oxygen

Formula Name in Aqueous Solution Name of Gaseous Species

HF hydrofluoric acid hydrogen fluoride

HCl hydrochloric acid hydrogen chloride

HBr hydrobromic acid hydrogen bromide

HI hydroiodic acid hydrogen iodide

HCN hydrocyanic acid hydrogen cyanide

H2S hydrosulfuric acid hydrogen sulfide

If an acid contains one or more H+ ions attached to oxygen, it is a derivative of oneof the common oxoanions, such as sulfate (SO4

2−) or nitrate (NO3−). These acids

contain as many H+ ions as are necessary to balance the negative charge on theanion, resulting in a neutral species such as H2SO4 and HNO3.

The names of acids are derived from the names of anions according to the followingrules:

1. If the name of the anion ends in -ate, then the name of the acidends in -ic. For example, because NO3

− is the nitrate ion, HNO3 is nitric

acid. Similarly, ClO4− is the perchlorate ion, so HClO4 is perchloric acid.

Two important acids are sulfuric acid (H2SO4) from the sulfate ion

(SO42−) and phosphoric acid (H3PO4) from the phosphate ion (PO4

3−).

These two names use a slight variant of the root of the anion name:sulfate becomes sulfuric and phosphate becomes phosphoric.

2. If the name of the anion ends in -ite, then the name of the acidends in -ous. For example, OCl− is the hypochlorite ion, and HOCl ishypochlorous acid; NO2

− is the nitrite ion, and HNO2 is nitrous acid;

and SO32− is the sulfite ion, and H2SO3 is sulfurous acid. The same roots


2.5 Acids and Bases 212

are used whether the acid name ends in -ic or -ous; thus, sulfite becomessulfurous.

The relationship between the names of the oxoacids and the parent oxoanions isillustrated in Figure 2.20 "The Relationship between the Names of the Oxoacids andthe Names of the Parent Oxoanions", and some common oxoacids are in Table 2.9"Some Common Oxoacids".

Figure 2.20 The Relationship between the Names of the Oxoacids and the Names of the Parent Oxoanions

Table 2.9 Some Common Oxoacids

Formula Name

HNO2 nitrous acid

HNO3 nitric acid

H2SO3 sulfurous acid

H2SO4 sulfuric acid

H3PO4 phosphoric acid

H2CO3 carbonic acid

HClO hypochlorous acid

HClO2 chlorous acid

HClO3 chloric acid

HClO4 perchloric acid



EXAMPLE 11

Name and give the formula for each acid.

a. the acid formed by adding a proton to the hypobromite ion (OBr−)b. the acid formed by adding two protons to the selenate ion (SeO4

2−)

Given: anion

Asked for: parent acid

Strategy:

Refer to Table 2.8 "Some Common Acids That Do Not Contain Oxygen" andTable 2.9 "Some Common Oxoacids" to find the name of the acid. If the acidis not listed, use the guidelines given previously.

Solution:

Neither species is listed in Table 2.8 "Some Common Acids That Do NotContain Oxygen" or Table 2.9 "Some Common Oxoacids", so we must use theinformation given previously to derive the name of the acid from the nameof the polyatomic anion.

a. The anion name, hypobromite, ends in -ite, so the name of the parent acidends in -ous. The acid is therefore hypobromous acid (HOBr).

b. Selenate ends in -ate, so the name of the parent acid ends in -ic. The acidis therefore selenic acid (H2SeO4).

Exercise

Name and give the formula for each acid.

a. the acid formed by adding a proton to the perbromate ion (BrO4−)

b. the acid formed by adding three protons to the arsenite ion (AsO33−)

Answer:

a. perbromic acid; HBrO4

b. arsenous acid; H3AsO3



Many organic compounds contain the carbonyl group38, in which there is acarbon–oxygen double bond. In carboxylic acids39, an –OH group is covalentlybonded to the carbon atom of the carbonyl group. Their general formula is RCO2H,

sometimes written as RCOOH:

where R can be an alkyl group, an aryl group, or a hydrogen atom. The simplestexample, HCO2H, is formic acid, so called because it is found in the secretions of

stinging ants (from the Latin formica, meaning “ant”). Another example is acetic acid(CH3CO2H), which is found in vinegar. Like many acids, carboxylic acids tend to

have sharp odors. For example, butyric acid (CH3CH2CH2CO2H), is responsible for

the smell of rancid butter, and the characteristic odor of sour milk and vomit is dueto lactic acid [CH3CH(OH)CO2H]. Some common carboxylic acids are shown in Figure

2.21 "Some Common Carboxylic Acids".

38. A carbon atom double-bondedto an oxygen atom. It is acharacteristic feature of manyorganic compounds, includingcarboxylic acids.

39. An organic compound thatcontains an −OH groupcovalently bonded to thecarbon atom of a carbonylgroup. The general formula ofa carboxylic acid is RCO2H.In water a carboxylic aciddissociates to produce an acidicsolution.



Figure 2.21 Some Common Carboxylic Acids

Although carboxylic acids are covalent compounds, when they dissolve in water,they dissociate to produce H+ ions (just like any other acid) and RCO2

− ions. Note

that only the hydrogen attached to the oxygen atom of the CO2 group dissociates to form an

H+ ion. In contrast, the hydrogen atom attached to the oxygen atom of an alcoholdoes not dissociate to form an H+ ion when an alcohol is dissolved in water. Thereasons for the difference in behavior between carboxylic acids and alcohols will bediscussed in Chapter 8 "Ionic versus Covalent Bonding".

Note the Pattern

Only the hydrogen attached to the oxygen atom of the CO2 group dissociates to

form an H+ ion.



Bases

We will present more comprehensive definitions of bases in later chapters, butvirtually every base you encounter in the meantime will be an ionic compound,such as sodium hydroxide (NaOH) and barium hydroxide [Ba(OH)2], that contain the

hydroxide ion and a metal cation. These have the general formula M(OH)n. It is

important to recognize that alcohols, with the general formula ROH, are covalentcompounds, not ionic compounds; consequently, they do not dissociate in water toform a basic solution (containing OH− ions). When a base reacts with any of theacids we have discussed, it accepts a proton (H+). For example, the hydroxide ion(OH−) accepts a proton to form H2O. Thus bases are also referred to as proton

acceptors.

Concentrated aqueous solutions of ammonia (NH3)

contain significant amounts of the hydroxide ion, eventhough the dissolved substance is not primarilyammonium hydroxide (NH4OH) as is often stated on the

label. Thus aqueous ammonia solution is also a commonbase. Replacing a hydrogen atom of NH3 with an alkyl

group results in an amine40 (RNH2), which is also a base.

Amines have pungent odors—for example, methylamine(CH3NH2) is one of the compounds responsible for the

foul odor associated with spoiled fish. The physiologicalimportance of amines is suggested in the word vitamin,which is derived from the phrase vital amines. The wordwas coined to describe dietary substances that were effective at preventing scurvy,rickets, and other diseases because these substances were assumed to be amines.Subsequently, some vitamins have indeed been confirmed to be amines.

Note the Pattern

Metal hydroxides (MOH) yield OH− ions and are bases, alcohols (ROH) do notyield OH− or H+ ions and are neutral, and carboxylic acids (RCO2H) yield H+ ions

and are acids.

40. An organic compound that hasthe general formula RNH2 ,where R is an alkyl group.Amines, like ammonia, arebases.



Summary

Common acids and the polyatomic anions derived from them have their ownnames and rules for nomenclature. The nomenclature of acids differentiatesbetween oxoacids, in which the H+ ion is attached to an oxygen atom of apolyatomic ion, and acids in which the H+ ion is attached to another element.Carboxylic acids are an important class of organic acids. Ammonia is animportant base, as are its organic derivatives, the amines.

KEY TAKEAWAY

• Common acids and polyatomic anions derived from them have their ownnames and rules for nomenclature.



CONCEPTUAL PROBLEMS

1. Name each acid.

a. HClb. HBrO3c. HNO3d. H2SO4e. HIO3

2. Name each acid.

a. HBrb. H2SO3c. HClO3d. HCNe. H3PO4

3. Name the aqueous acid that corresponds to each gaseous species.

a. hydrogen bromideb. hydrogen cyanidec. hydrogen iodide


a.

b.


a.



b.

6. When each compound is added to water, is the resulting solution acidic,neutral, or basic?

a. CH3CH2OHb. Mg(OH)2c. C6H5CO2Hd. LiOHe. C3H7CO2Hf. H2SO4

7. Draw the structure of the simplest example of each type of compound.

a. alkaneb. alkenec. alkyned. aromatic hydrocarbone. alcoholf. carboxylic acidg. amineh. cycloalkane

8. Identify the class of organic compound represented by each compound.

a.

b. CH3CH2OH

c. HC≡CH

d.

e. C3H7NH2



f. CH3CH=CHCH2CH3

g.

h.

9. Identify the class of organic compound represented by each compound.

a.

b.

c.

d.

e.

f. CH3C≡CH

g.

h.



NUMERICAL PROBLEMS


a. hypochlorous acidb. perbromic acidc. hydrobromic acidd. sulfurous acide. sodium perbromate


a. hydroiodic acidb. hydrogen sulfidec. phosphorous acidd. perchloric acide. calcium hypobromite


a. HBrb. H2SO3c. HCNd. HClO4e. NaHSO4


a. H2SO4b. HNO2c. K2HPO4d. H3PO3e. Ca(H2PO4)2·H2O



2.6 Industrially Important Chemicals

LEARNING OBJECTIVE

1. To appreciate the scope of the chemical industry and its contributions tomodern society.

It isn’t easy to comprehend the scale on which the chemical industry must operateto supply the huge amounts of chemicals required in modern industrial societies.Figure 2.22 "Top 25 Chemicals Produced in the United States in 2002*" lists thenames and formulas of the chemical industry’s “top 25” for 2002—the 25 chemicalsproduced in the largest quantity in the United States that year—along with theamounts produced, in billions of pounds. To put these numbers in perspective,consider that the 88.80 billion pounds of sulfuric acid produced in the United Statesin 2002 has a volume of 21.90 million cubic meters (2.19 × 107 m3), enough to fill thePentagon, probably the largest office building in the world, about 22 times.

Figure 2.22 Top 25 Chemicals Produced in the United States in 2002*


223

According to Figure 2.22 "Top 25 Chemicals Produced in the United States in 2002*",11 of the top 15 compounds produced in the United States are inorganic, and thetotal mass of inorganic chemicals produced is almost twice the mass of organicchemicals. Yet the diversity of organic compounds used in industry is such thatover half of the top 25 compounds (13 out of 25) are organic.

Why are such huge quantities of chemical compounds produced annually? They areused both directly as components of compounds and materials that we encounteron an almost daily basis and indirectly in the production of those compounds andmaterials. The single largest use of industrial chemicals is in the production offoods: 7 of the top 15 chemicals are either fertilizers (ammonia, urea, andammonium nitrate) or used primarily in the production of fertilizers (sulfuric acid,nitric acid, nitrogen, and phosphoric acid). Many of the organic chemicals on thelist are used primarily as ingredients in the plastics and related materials that areso prevalent in contemporary society. Ethylene and propylene, for example, areused to produce polyethylene and polypropylene, which are made into plastic milkbottles, sandwich bags, indoor-outdoor carpets, and other common items. Vinylchloride, in the form of polyvinylchloride, is used in everything from pipes to floortiles to trash bags. Though not listed in Figure 2.22 "Top 25 Chemicals Produced inthe United States in 2002*", butadiene and carbon black are used in themanufacture of synthetic rubber for tires, and phenol and formaldehyde areingredients in plywood, fiberglass, and many hard plastic items.

We do not have the space in this text to consider the applications of all thesecompounds in any detail, but we will return to many of them after we havedeveloped the concepts necessary to understand their underlying chemistry.Instead, we conclude this chapter with a brief discussion of petroleum refining as itrelates to gasoline and octane ratings and a look at the production and use of thetopmost industrial chemical, sulfuric acid.

Petroleum

The petroleum that is pumped out of the ground at locations around the world is acomplex mixture of several thousand organic compounds, including straight-chainalkanes, cycloalkanes, alkenes, and aromatic hydrocarbons with four to severalhundred carbon atoms. The identities and relative abundances of the componentsvary depending on the source. So Texas crude oil is somewhat different from SaudiArabian crude oil. In fact, the analysis of petroleum from different deposits canproduce a “fingerprint” of each, which is useful in tracking down the sources ofspilled crude oil. For example, Texas crude oil is “sweet,” meaning that it contains asmall amount of sulfur-containing molecules, whereas Saudi Arabian crude oil is“sour,” meaning that it contains a relatively large amount of sulfur-containingmolecules.


2.6 Industrially Important Chemicals 224

Gasoline

Petroleum is converted to useful products such as gasoline in three steps:distillation, cracking, and reforming. Recall from Chapter 1 "Introduction toChemistry" that distillation separates compounds on the basis of their relativevolatility, which is usually inversely proportional to their boiling points. Part (a) inFigure 2.23 "The Distillation of Petroleum" shows a cutaway drawing of a columnused in the petroleum industry for separating the components of crude oil. Thepetroleum is heated to approximately 400°C (750°F), at which temperature it hasbecome a mixture of liquid and vapor. This mixture, called the feedstock, isintroduced into the refining tower. The most volatile components (those with thelowest boiling points) condense at the top of the column where it is cooler, whilethe less volatile components condense nearer the bottom. Some materials are sononvolatile that they collect at the bottom without evaporating at all. Thus thecomposition of the liquid condensing at each level is different. These differentfractions, each of which usually consists of a mixture of compounds with similarnumbers of carbon atoms, are drawn off separately. Part (b) in Figure 2.23 "TheDistillation of Petroleum" shows the typical fractions collected at refineries, thenumber of carbon atoms they contain, their boiling points, and their ultimate uses.These products range from gases used in natural and bottled gas to liquids used infuels and lubricants to gummy solids used as tar on roads and roofs.

Figure 2.23 The Distillation of Petroleum

(a) This is a diagram of a distillation column used for separating petroleum fractions. (b) Petroleum fractionscondense at different temperatures, depending on the number of carbon atoms in the molecules, and are drawn offfrom the column. The most volatile components (those with the lowest boiling points) condense at the top of thecolumn, and the least volatile (those with the highest boiling points) condense at the bottom.



The economics of petroleum refining are complex. For example, the marketdemand for kerosene and lubricants is much lower than the demand for gasoline,yet all three fractions are obtained from the distillation column in comparableamounts. Furthermore, most gasolines and jet fuels are blends with very carefullycontrolled compositions that cannot vary as their original feedstocks did. To makepetroleum refining more profitable, the less volatile, lower-value fractions must beconverted to more volatile, higher-value mixtures that have carefully controlledformulas. The first process used to accomplish this transformation is cracking41, inwhich the larger and heavier hydrocarbons in the kerosene and higher-boiling-point fractions are heated to temperatures as high as 900°C. High-temperaturereactions cause the carbon–carbon bonds to break, which converts the compoundsto lighter molecules similar to those in the gasoline fraction. Thus in cracking, astraight-chain alkane with a number of carbon atoms corresponding to thekerosene fraction is converted to a mixture of hydrocarbons with a number ofcarbon atoms corresponding to the lighter gasoline fraction. The second processused to increase the amount of valuable products is called reforming42; it is thechemical conversion of straight-chain alkanes to either branched-chain alkanes ormixtures of aromatic hydrocarbons. Using metals such as platinum brings about thenecessary chemical reactions. The mixtures of products obtained from cracking andreforming are separated by fractional distillation.

Octane Ratings

The quality of a fuel is indicated by its octane rating43, which is a measure of itsability to burn in a combustion engine without knocking or pinging. Knocking andpinging signal premature combustion (Figure 2.24 "The Burning of Gasoline in anInternal Combustion Engine"), which can be caused either by an engine malfunctionor by a fuel that burns too fast. In either case, the gasoline-air mixture detonates atthe wrong point in the engine cycle, which reduces the power output and candamage valves, pistons, bearings, and other engine components. The variousgasoline formulations are designed to provide the mix of hydrocarbons least likelyto cause knocking or pinging in a given type of engine performing at a particularlevel.

41. A process in petroleumrefining in which the largerand heavier hydrocarbons inkerosene and higher-boiling-point fractions are heated tohigh temperatures, causing thecarbon–carbon bonds to break(“crack”), thus producing amore volatile mixture.

42. The second process used inpetroleum refining, which isthe chemical conversion ofstraight-chain alkanes to eitherbranched-chain alkanes ormixtures of aromatichydrocarbons.

43. A measure of a fuel’s ability toburn in a combustion enginewithout knocking or pinging(indications of prematurecombustion). The higher theoctane rating, the higherquality the fuel.



Figure 2.24 The Burning of Gasoline in an Internal Combustion Engine

(a) Normally, fuel is ignited by the spark plug, and combustion spreads uniformly outward. (b) Gasoline with anoctane rating that is too low for the engine can ignite prematurely, resulting in uneven burning that causesknocking and pinging.

The octane scale was established in 1927 using a standard test engine and two purecompounds: n-heptane and isooctane (2,2,4-trimethylpentane). n-Heptane, whichcauses a great deal of knocking on combustion, was assigned an octane rating of 0,whereas isooctane, a very smooth-burning fuel, was assigned an octane rating of100. Chemists assign octane ratings to different blends of gasoline by burning asample of each in a test engine and comparing the observed knocking with theamount of knocking caused by specific mixtures of n-heptane and isooctane. Forexample, the octane rating of a blend of 89% isooctane and 11% n-heptane is simplythe average of the octane ratings of the components weighted by the relativeamounts of each in the blend. Converting percentages to decimals, we obtain theoctane rating of the mixture:

0.89(100) + 0.11(0) = 89

A gasoline that performs at the same level as a blend of 89% isooctane and 11% n-heptane is assigned an octane rating of 89; this represents an intermediate grade ofgasoline. Regular gasoline typically has an octane rating of 87; premium has a ratingof 93 or higher.

As shown in Figure 2.25 "The Octane Ratings of Some Hydrocarbons and CommonAdditives", many compounds that are now available have octane ratings greaterthan 100, which means they are better fuels than pure isooctane. In addition,



antiknock agents, also called octane enhancers, have been developed. One of themost widely used for many years was tetraethyllead [(C2H5)4Pb], which at

approximately 3 g/gal gives a 10–15-point increase in octane rating. Since 1975,however, lead compounds have been phased out as gasoline additives because theyare highly toxic. Other enhancers, such as methyl t-butyl ether (MTBE), have beendeveloped to take their place. They combine a high octane rating with minimalcorrosion to engine and fuel system parts. Unfortunately, when gasoline containingMTBE leaks from underground storage tanks, the result has been contamination ofthe groundwater in some locations, resulting in limitations or outright bans on theuse of MTBE in certain areas. As a result, the use of alternative octane enhancerssuch as ethanol, which can be obtained from renewable resources such as corn,sugar cane, and, eventually, corn stalks and grasses, is increasing.

Figure 2.25 The Octane Ratings of Some Hydrocarbons and Common Additives



EXAMPLE 12

You have a crude (i.e., unprocessed or straight-run) petroleum distillateconsisting of 10% n-heptane, 10% n-hexane, and 80% n-pentane by mass,with an octane rating of 52. What percentage of MTBE by mass would youneed to increase the octane rating of the distillate to that of regular-gradegasoline (a rating of 87), assuming that the octane rating is directlyproportional to the amounts of the compounds present? Use theinformation presented in Figure 2.25 "The Octane Ratings of SomeHydrocarbons and Common Additives".

Given: composition of petroleum distillate, initial octane rating, and finaloctane rating

Asked for: percentage of MTBE by mass in final mixture

Strategy:

A Define the unknown as the percentage of MTBE in the final mixture. Thensubtract this unknown from 100% to obtain the percentage of petroleumdistillate.

B Multiply the percentage of MTBE and the percentage of petroleumdistillate by their respective octane ratings; add these values to obtain theoverall octane rating of the new mixture.

C Solve for the unknown to obtain the percentage of MTBE needed.

Solution:

A The question asks what percentage of MTBE will give an overall octanerating of 87 when mixed with the straight-run fraction. From Figure 2.25"The Octane Ratings of Some Hydrocarbons and Common Additives", theoctane rating of MTBE is 116. Let x be the percentage of MTBE, and let 100 − xbe the percentage of petroleum distillate.

B Multiplying the percentage of each component by its respective octanerating and setting the sum equal to the desired octane rating of the mixture(87) times 100 gives



C Solving the equation gives x = 55%. Thus the final mixture must contain55% MTBE by mass.

To obtain a composition of 55% MTBE by mass, you would have to add morethan an equal mass of MTBE (actually 0.55/0.45, or 1.2 times) to the straight-run fraction. This is 1.2 tons of MTBE per ton of straight-run gasoline, whichwould be prohibitively expensive. Thus there are sound economic reasonsfor reforming the kerosene fractions to produce toluene and other aromaticcompounds, which have high octane ratings and are much cheaper thanMTBE.

Exercise

As shown in Figure 2.25 "The Octane Ratings of Some Hydrocarbons andCommon Additives", toluene is one of the fuels suitable for use inautomobile engines. How much toluene would have to be added to a blend ofthe petroleum fraction in this example containing 15% MTBE by mass toincrease the octane rating to that of premium gasoline (93)?

Answer: The final blend is 56% toluene by mass, which requires a ratio of56/44, or 1.3 tons of toluene per ton of blend.

Sulfuric Acid

Sulfuric acid is one of the oldest chemical compounds known. It was probably firstprepared by alchemists who burned sulfate salts such as FeSO4·7H2O, called green

vitriol from its color and glassy appearance (from the Latin vitrum, meaning“glass”). Because pure sulfuric acid was found to be useful for dyeing textiles,enterprising individuals looked for ways to improve its production. By the mid-18thcentury, sulfuric acid was being produced in multiton quantities by the lead-chamberprocess, which was invented by John Roebuck in 1746. In this process, sulfur wasburned in a large room lined with lead, and the resulting fumes were absorbed inwater.

final octane rating of mixture = 87(100)= 52(100 − x) + 116x

= 5200 − 52x + 116x

= 5200 + 64x



Transporting sulfur. A traincarries elemental sulfur throughthe White Canyon of theThompson River in BritishColumbia, Canada.

Production

The production of sulfuric acid today is likely to start with elemental sulfurobtained through an ingenious technique called the Frasch process, which takesadvantage of the low melting point of elemental sulfur (115.2°C). Large deposits ofelemental sulfur are found in porous limestone rocks in the same geologicalformations that often contain petroleum. In the Frasch process, water at hightemperature (160°C) and high pressure is pumped underground to melt the sulfur,and compressed air is used to force the liquid sulfur-water mixture to the surface(Figure 2.26 "Extraction of Elemental Sulfur from Underground Deposits"). Thematerial that emerges from the ground is more than 99% pure sulfur. After itsolidifies, it is pulverized and shipped in railroad cars to the plants that producesulfuric acid, as shown here.

Figure 2.26 Extraction of Elemental Sulfur from Underground Deposits



In the Frasch process for extracting sulfur, very hot water at high pressure is injected into the sulfur-containingrock layer to melt the sulfur. The resulting mixture of liquid sulfur and hot water is forced up to the surface bycompressed air.

An increasing number of sulfuric acid manufacturers have begun to use sulfurdioxide (SO2) as a starting material instead of elemental sulfur. Sulfur dioxide is

recovered from the burning of oil and gas, which contain small amounts of sulfurcompounds. When not recovered, SO2 is released into the atmosphere, where it is

converted to an environmentally hazardous form that leads to acid rain (Chapter 4"Reactions in Aqueous Solution").

If sulfur is the starting material, the first step in the production of sulfuric acid isthe combustion of sulfur with oxygen to produce SO2. Next, SO2 is converted to SO3

by the contact process, in which SO2 and O2 react in the presence of V2O5 to achieve

about 97% conversion to SO3. The SO3 can then be treated with a small amount of

water to produce sulfuric acid. Usually, however, the SO3 is absorbed in

concentrated sulfuric acid to produce oleum, a more potent form called fumingsulfuric acid. Because of its high SO3 content (approximately 99% by mass), oleum is

cheaper to ship than concentrated sulfuric acid. At the point of use, the oleum isdiluted with water to give concentrated sulfuric acid (very carefully because dilutiongenerates enormous amounts of heat). Because SO2 is a pollutant, the small

amounts of unconverted SO2 are recovered and recycled to minimize the amount

released into the air.

Uses

Two-thirds of the sulfuric acid produced in the United States is used to makefertilizers, most of which contain nitrogen, phosphorus, and potassium (in a formcalled potash). In earlier days, phosphate-containing rocks were simply ground upand spread on fields as fertilizer, but the extreme insolubility of many salts thatcontain the phosphate ion (PO4

3−) limits the availability of phosphorus from these

sources. Sulfuric acid serves as a source of protons (H+ ions) that react withphosphate minerals to produce more soluble salts containing HPO4

2− or H2PO4− as

the anion, which are much more readily taken up by plants. In this context, sulfuricacid is used in two principal ways: (1) the phosphate rocks are treated withconcentrated sulfuric acid to produce “superphosphate,” a mixture of 32% CaHPO4

and Ca(H2PO4)2·H2O, 50% CaSO4·2H2O, approximately 3% absorbed phosphoric acid,

and other nutrients; and (2) sulfuric acid is used to produce phosphoric acid(H3PO4), which can then be used to convert phosphate rocks to “triple

superphosphate,” which is largely Ca(H2PO4)2·H2O.



Sulfuric acid is also used to produce potash, one of the other major ingredients infertilizers. The name potash originally referred to potassium carbonate (obtained byboiling wood ashes with water in iron pots), but today it also refers to compoundssuch as potassium hydroxide (KOH) and potassium oxide (K2O). The usual source of

potassium in fertilizers is actually potassium sulfate (K2SO4), which is produced by

several routes, including the reaction of concentrated sulfuric acid with solidpotassium chloride (KCl), which is obtained as the pure salt from mineral deposits.

Summary

Many chemical compounds are prepared industrially in huge quantities andused to produce foods, fuels, plastics, and other such materials. Petroleumrefining takes a complex mixture of naturally occurring hydrocarbons as afeedstock and, through a series of steps involving distillation, cracking, andreforming, converts them to mixtures of simpler organic compounds withdesirable properties. A major use of petroleum is in the production of motorfuels such as gasoline. The performance of such fuels in engines is described bytheir octane rating, which depends on the identity of the compounds presentand their relative abundance in the blend.

Sulfuric acid is the compound produced in the largest quantity in the industrialworld. Much of the sulfur used in the production of sulfuric acid is obtained viathe Frasch process, in which very hot water forces liquid sulfur out of theground in nearly pure form. Sulfuric acid is produced by the reaction of sulfurdioxide with oxygen in the presence of vanadium(V) oxide (the contactprocess), followed by absorption of the sulfur trioxide in concentrated sulfuricacid to produce oleum. Most sulfuric acid is used to prepare fertilizers.

KEY TAKEAWAY

• Many chemical compounds are prepared industrially in huge quantitiesto prepare the materials we need and use in our daily lives.



CONCEPTUAL PROBLEMS

1. Describe the processes used for converting crude oil to transportation fuels.

2. If your automobile engine is knocking, is the octane rating of your gasoline toolow or too high? Explain your answer.

3. Tetraethyllead is no longer used as a fuel additive to prevent knocking.Instead, fuel is now marketed as “unleaded.” Why is tetraethyllead no longerused?

4. If you were to try to extract sulfur from an underground source, what processwould you use? Describe briefly the essential features of this process.

5. Why are phosphate-containing minerals used in fertilizers treated withsulfuric acid?

ANSWER

5. Phosphate salts contain the highly-charged PO43− ion, salts of which are often

insoluble. Protonation of the PO43− ion by strong acids such as H2SO4 leads to

the formation of the HPO42− and H2PO4

− ions. Because of their decreasednegative charge, salts containing these anions are usually much more soluble,allowing the anions to be readily taken up by plants when they are applied asfertilizer.

NUMERICAL PROBLEM

1. In Example 12, the crude petroleum had an overall octane rating of 52. What isthe composition of a solution of MTBE and n-heptane that has this octanerating?



2.7 End-of-Chapter Material


235

APPLICATION PROBLEMS

Problems marked with a ♦ involve multiple concepts.

1. Carbon tetrachloride (CCl4) was used as a dry cleaning solvent until it wasfound to cause liver cancer. Based on the structure of chloroform given inSection 2.1 "Chemical Compounds", draw the structure of carbontetrachloride.

2. Ammonium nitrate and ammonium sulfate are used in fertilizers as a source ofnitrogen. The ammonium cation is tetrahedral. Refer to Section 2.1 "ChemicalCompounds" to draw the structure of the ammonium ion.

3. The white light in fireworks displays is produced by burning magnesium in air,which contains oxygen. What compound is formed?

4. Sodium hydrogen sulfite, which is used for bleaching and swelling leather andto preserve flavor in almost all commercial wines, is made from sulfur dioxide.What are the formulas for these two sulfur-containing compounds?

5. Carbonic acid is used in carbonated drinks. When combined with lithiumhydroxide, it produces lithium carbonate, a compound used to increase thebrightness of pottery glazes and as a primary treatment for depression andbipolar disorder. Write the formula for both of these carbon-containingcompounds.

6. Vinegar is a dilute solution of acetic acid, an organic acid, in water. Whatgrouping of atoms would you expect to find in the structural formula for aceticacid?

7. ♦ Sodamide, or sodium amide, is prepared from sodium metal and gaseousammonia. Sodamide contains the amide ion (NH2

−), which reacts with water toform the hydroxide anion by removing an H+ ion from water. Sodium amide isalso used to prepare sodium cyanide.

a. Write the formula for each of these sodium-containing compounds.b. What are the products of the reaction of sodamide with water?

8. A mixture of isooctane, n-pentane, and n-heptane is known to have an octanerating of 87. Use the data in Figure 2.25 "The Octane Ratings of SomeHydrocarbons and Common Additives" to calculate how much isooctane and n-heptane are present if the mixture is known to contain 30% n-pentane.

9. A crude petroleum distillate consists of 60% n-pentane, 25% methanol, and theremainder n-hexane by mass (Figure 2.25 "The Octane Ratings of SomeHydrocarbons and Common Additives").


2.7 End-of-Chapter Material 236

a. What is the octane rating?b. How much MTBE would have to be added to increase the octane rating to

93?

10. Premium gasoline sold in much of the central United States has an octanerating of 93 and contains 10% ethanol. What is the octane rating of the gasolinefraction before ethanol is added? (See Figure 2.25 "The Octane Ratings of SomeHydrocarbons and Common Additives".)

ANSWERS

1.

3. MgO, magnesium oxide

5. Carbonic acid is H2CO3; lithium carbonate is Li2CO3.

7. a. Sodamide is NaNH2, and sodium cyanide is NaCN.b. Sodium hydroxide (NaOH) and ammonia (NH3).

9. a. 68b. 52 g of MTBE must be added to 48 g of the crude distillate.


2.7 End-of-Chapter Material 237

molecules, ions, and chemical formulas · the molecule has 4 phosphorus atoms and 3 sulfur atoms....

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